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Platinum as a HOI/I Redox Electrode and its Mixed Potential in the Oscillatory Briggs-Rauscher Reaction Gábor Holló, Kristóf Kály-Kullai, Thuy Bich Lawson, Zoltan Noszticzius, Maria Wittmann, Norbert Thomas Muntean, Stanley D. Furrow, and Guy E. Schmitz J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/acs.jpca.6b10243 • Publication Date (Web): 15 Dec 2016 Downloaded from http://pubs.acs.org on December 19, 2016

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The Journal of Physical Chemistry

Platinum as a HOI/I2 Redox Electrode and its Mixed Potential in the Oscillatory Briggs-Rauscher Reaction Gábor Hollóa, Kristóf Kály-Kullaia, Thuy B. Lawsona, Zoltán Noszticziusa, Maria Wittmanna, Norbert Munteana,b, Stanley D. Furrowc, Guy Schmitzd a

Department of Physics, Budapest University of Technology and Economics,

H-1521 Budapest, Hungary b

Department of Physical Chemistry, Babes-Bolyai University,

RO-400028 Cluj-Napoca, Romania c

Penn State Berks College, The Pennsylvania State University, Reading,

Pennsylvania 19610 USA d

Faculty of Applied Sciences, Université Libre de Bruxelles,

CP165/63, Av. F. Roosevelt 50, 1050 Brussels, Belgium

Abstract. Pt is a common redox electrode used to follow oscillations qualitatively in the BriggsRauscher (BR) and the Bray-Liebhafsky (BL) reactions from the time of their discovery. While 1 ACS Paragon Plus Environment


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the potential oscillations of the electrode reflect the temporal pattern of the reaction properly, there is no general agreement how that potential is determined by the components of the reaction mixture. In this paper first we investigate how iodine species in different oxidation states affect the potential of a Pt electrode. It was found that I(+3) and I(+5) species do not affect the potential; only I–, I2 and HOI may have an influence. While the latter three species are always present simultaneously as participants of the rapid iodine hydrolysis equilibrium, it was found that below and above the so-called hydrolysis limit potential (HLP; where the iodide and HOI concentrations are equal) the actual potential determining redox couple is different. Below the HLP it is the traditional I2/I– redox couple but above the HLP it is the HOI/I2 redox pair which determines the potential of a Pt electrode. That change in the potential control mechanism was proved experimentally by exchange current measurements. In addition, from the potential response of the Pt electrode below and above the HLP, it was possible to calculate the equilibrium constant of the iodine hydrolysis as K°H = (4.97 ± 0.20)10–13 M2 in rather good agreement with earlier measurements. We also studied the perturbing effect of H2O2 on the previously mentioned potentials. The concentration of H2O2 was 0.66 M, as in the BR reaction studied here. It was found that below the HLP, the perturbing effect of H2O2 was minimal, but above the HLP, H2O2 shifted the mixed potential considerably down toward the HLP. In our experiments with the BR reaction the potential oscillations of the Pt electrode crossed the HLP, indicating that from time to time HOI concentration exceeds that of the iodide. We can conclude that while the perturbing effect of H2O2 prevents calculating concentrations from Pt potentials above the HLP, [I–]/[I2]1/2 ratios can be calculated as a good approximation from Pt potentials below the HLP.

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1. Introduction The classical Briggs-Rauscher (BR) oscillating reaction is the reaction between H2O2, malonic acid and iodate in the presence of Mn(II) catalyst in an acidic solution1. Platinum and AgI electrodes were used to follow oscillations from the early beginning2. In most cases, however, these electrode potential vs. time diagrams (“potentiometric traces”) are used only semi-quantitatively. For example, the amplitude of the Pt potential oscillations is usually not taken into account when a BR oscillator is applied for analytical purposes3. In that case, only the time period of the reaction or the number of oscillations is regarded. Moreover, there is no general agreement concerning the potential determining redox pair for a Pt electrode in the BR or in a BL reaction. Some authors assume that it is the HO2/H2O2 in a BR reaction4 while others suggest that the I2/I– redox couple plays the role of the potential determining redox pair in the BL system5. There were somewhat similar uncertainties regarding the potential of a AgI electrode. Especially it was not clear how hypoiodous acid might affect the potential of a AgI or a Pt electrode. Thus, we have started a systematic research with the aim to describe how the various intermediates of the BR reaction affect the potential of these two most frequently used indicator electrodes. In two previous publications6,7 we dealt with the potential response of the AgI electrode and now we focus on the platinum electrode. As we should like to compare the behavior of the two electrodes, to show both the similarities and also the dissimilarities in their potential response, first we summarize briefly our previous results with the AgI electrode and then we give an overview of our present experiments with the Pt electrode. 1.1. Potential response of the AgI iodide selective electrode for oxyiodine species. (The essence of previous results important for this work.)



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In most studies, it is assumed that the potential of a AgI electrode is determined always by the iodide concentration of the bulk via the Nernst equation. In two recent papers6,7 we have shown, however, that this is not true when the electrode potential is above the so-called solubility limit potential or SLP. (The SLP can be measured when the silver and iodide ion concentrations are equal at the surface of the electrode: [Ag+]SL = [I–]SL = “solubility limit” concentrations and

s

s;

here the index “SL” refers to the

is the solubility product of AgI.) An equilibrium potential

above the SLP can be measured in an electrolyte mixture only if the bulk contains silver ions in concentrations higher than

s.

Most chemical oscillators do not contain silver ions, however.

Consequently, any silver ions appearing in such systems should originate from the corrosion of the AgI electrode by HOI. This way, above the SLP and in the absence of bulk silver ions, the iodide selective electrode actually works as a HOI selective electrode. It was shown6 that bulk iodine might also influence the measured potential. However, that effect is a few mV only as long as [I2]B stays below 10 M (see Fig. 5a in Ref. 6). In another previous publication7, we also tested the effect of other oxyiodine species, like iodous and iodic acids, on the electrode potential. In theory, these acids could contribute to the corrosion of AgI and the extra silver ions produced this way might affect the electrode potential. It was found, however, that in the presence of HOI the contribution of these species to the corrosion potential is below the detection limit. 1.2. Potential response of the platinum electrode for oxyiodine species: research aims of the present study. The main aim of the present work is to study the potential response of the Pt electrode for oxyiodine species, especially for HOI, because HOI can appear periodically in the BR reaction6. It is well known that in the presence of I2 and I–, the potential of the Pt electrode Pt at 25 °C can 4 ACS Paragon Plus Environment

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be written8,9 as Pt = o t + 59.16 mVlg

2 –

.

However, when HOI is present in a high concentration in a mixture of I2, I– and HOI, the rapid iodine hydrolysis equilibrium drives [I–] to a very low value. In this case, while the above formula is still valid from a thermodynamical point of view, it is questionable whether a stable electrode potential can be measured when the potential determining iodide ion concentration is so low. To test the Pt electrode in this respect first we decided to titrate a HOI solution with I– in the simultaneous presence of Pt and AgI indicator electrodes (see the apparatus in Fig. 1. and the results in Fig. 2, both in the Experimental part). Before the equivalence point [I–] is low because of the high [HOI] thus a stable electrode potential in this region would suggest that in this case, Pt potentials are not controlled directly by the small I– concentration and thus by the I2/I– redox couple, but rather by the HOI/I2 redox couple. Obviously there is no distinction in the calculated potential, as the two couples participate in the same hydrolysis equilibrium: –

I2 + H2O ⇌ HOI + I– + H+

2

so the electrode potential can be written in an alternative form: Pt = o t + 59.16 mVlg

2 –

= o t + 59.16 mVlg

+ 59.16 mVlg

2

= o t + 59.16 mV lg

At the equivalence point, when [HOI]HL = [I–]HL =

2

L

(the subscript L refers to the “hydrolysis limit”) a hydrolysis limit potential (HLP) of the Pt electrode Pt(HL) can be defined as Pt(HL) = 0 t + 59.16 mVlg 5 ACS Paragon Plus Environment

.

2

.


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As we will see, HLP plays a somewhat analogous role for the Pt electrode as SLP for the AgI electrode. In case of obtaining stable over-HLP Pt electrode potentials in the presence of HOI we also aimed to measure the Pt vs. lg

2

calibration line (this experiment will be shown in Fig. 3).

It can be seen from the two alternative forms of the Pt potential that increasing the iodine concentration should have a different effect below and above the HLP. Below the HLP I2 is the oxidized form in the I2/I– redox couple, thus increasing its concentration increases the Pt potential. However, above the HLP I2 is the reduced form in the HOI/I2 redox couple, thus increasing its concentration should decrease the Pt potential. The effect of adding iodine below and above the HLP will be presented in Fig. 4. A more direct possibility to prove that below and above the HLP different redox couples control the Pt potential is to perform exchange current measurements. In the course of a titration of HOI by I– in the presence of a relatively high I2 concentration the I– concentration increases and the HOI concentration decreases monotonically, while the I2 concentration is practically constant (its increase is relatively small). Thus if the exchange current and the Pt potential were determined by the concentration of the I– and I2 species then the exchange current should increase monotonically during the whole titration. However, supposing that above the HLP the actual potential controlling species are HOI and I2 then the exchange current should decrease monotonically reflecting the decreasing HOI concentration until the HLP, below which I– and I2 species take over the control. Thus supposing that below and above the HLP the actual potential controlling redox couples are not the same, the exchange current is not a monotonic function any more but it should have a minimum at the HLP. The method for determining the exchange current will be demonstrated in Fig. 5, and our experimental result will be shown in Fig. 6. 6 ACS Paragon Plus Environment

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We also aimed to test the effect of iodous and iodic acids on the Pt potential because in theory I(+3) and I(+5) could also affect the redox potential if their reduction by electrons could occur with a measurable rate on the Pt surface. Fig. 7 displays this experiment. We remark that it is an important difference between the Pt and the AgI electrodes that Pt cannot be corroded. Thus, in the absence of any surface reactions, there is equilibrium, and the surface and bulk HOI concentrations are equal. 1.3. Perturbation of the Pt and AgI electrode potentials by H2O2. In BR systems iodine occurs in five different oxidation states: I(–1), I(0), I(+1), I(+3), and I(+5), but the potential of the AgI electrode is not perturbed by I(+3) or I(+5) iodine species, and – as it will be shown in this paper – neither is that of the Pt electrode. However, in BR systems there is an additional component, namely H2O2, which can strongly affect at least the potential of the platinum electrode10. To check the perturbing effect of H2O2 we decided to measure potentials of two Pt electrodes, and for comparison also that of two AgI electrodes. The “test” electrode is to be submerged into a 25 mM H2SO4 solution with oxyiodine species while the “perturbing” electrode into a solution containing 0.66 M H2O2 (as in the BR medium applied here) and 25 mM H2SO4. The two solutions are to be connected with a salt bridge as shown in Fig. 8. Mixed potential11 is measured when there are more than one redox couples that determine the potential of the electrode. In the classical arrangement a single electrode is submerged in a solution containing different redox couples. The resulting potential is a non-equilibrium one, it is somewhere between the individual potentials (a kind of weighted average). No resulting current is flowing through the electrode in this case, however. A similar, but not identical situation is when two different electrodes are submerged into



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the same electrolyte and they are connected12. While each electrode could measure an equilibrium potential individually, this is not the case when they are connected: such a “combined” electrode exhibits a non-equilibrium potential which is again somewhere between the individual equilibrium potentials, and a macroscopic current is flowing through the electrodes. This analogous situation is also called a “mixed potential”12. Our case is a combination of the above situations: the classical arrangement does not permit to measure a mixed potential when the redox couples can react with each other, so we separated the redox couples to different solutions and we duplicated the Pt electrodes (we have two identical electrodes in touch with different redox couples). The situation is analogous to the previous ones with a current flowing through the electrodes and we call it also “mixed potentials”. The result of this experiment is shown in Fig. 9. 1.4. Simultaneous potential oscillations of Pt and AgI electrodes in a BR system. Finally the potentials of Pt and AgI electrodes will be recorded simultaneously in an oscillatory BR reaction (Fig. 10). The aim of these experiments is to observe whether or not the potential oscillations cross their limit values (i.e. HLP for Pt and SLP for AgI).

2. Experimental 2.1. Chemicals. KIO3 (Sigma Aldrich, ACS reagent,

.5 ),

( iedelde a n, puriss.

pa), H2SO4 (96%, Thomasker, pa), AgNO3 (Reanal, pa), I2 (Reanal, pa), MnSO4 4

2O

(Reanal,

pa), H2O2 (Fluka, puriss. pa ACS; ≥ 30 ), malonic acid (Fluka, puriss.), dimethyl sulfoxide (Sigma Aldrich), and dichloromethane (Sigma Aldrich) were used as received. All solutions were prepared with doubly distilled water. 2.2. Preparation of the solutions used in the experiments. 8 ACS Paragon Plus Environment

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2.2.1. Preparation of I(+1) solutions in concentrated sulphuric acid and in dichloromethane (DCM). I(+1) solutions in H2SO4 were prepared in a comproportionation reaction between I(0) and I(+5) by dissolving I2 and KIO3 in concentrated (96%) sulphuric acid. These solutions were extracted with DCM to yield I(+1) solutions in the organic solvent. The recipe of these preparations was the same as described previously6. The “I(+1) in DCM” solutions could be stored for weeks at 5 °C and were used to prepare HOI solutions freshly before the experiments. 2.2.2. Preparation of HOI solutions in 25 mM and in 1 M sulphuric acid. Dilute aqueous HOI solutions in 25 mM or 1 M sulphuric acid were prepared by adding 50-200 L of I(+1) in DCM stock solution to 50 mL of 25 mM or 1 M H2SO4 under intense stirring for about 1 minute. The concentration of HOI was determined by titration. Preparation of I(+3) solution in concentrated sulphuric acid. I(+3) solution was prepared also with comproportion as in the case of I(+1) but applying a higher amount of iodate, in the reaction between KIO3 dissolved in concentrated H2SO4 and I2 dissolved in DCM. The exact recipe can be found in Ref. 7 as I(+3)–C. 2.2.3. Preparation of dilute iodine solutions. When preparing these solutions one has to consider that a “pure” iodine solution always contains some iodide (which might be a contamination already present in the iodine crystals but produced mostly in these solutions by the slow disproportionation of iodine to iodide and iodate). This small amount of iodide could not be neglected in our case, especially in the sensitive calibration measurements with the Pt electrode. So before using these stoichiometric solutions we eliminated most of the contaminating iodide by adding some AgNO3 as described below (see “Calibration by adding solutions”).


–

or Ag+ ions to I2


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2.2.4. 0.1 M iodine stock solution in dimethyl sulfoxide (DMSO). The appropriate amount (253.8 mg) of iodine crystals was added to 10 mL DMSO and mixed for 5-10 minutes. The iodide concentration of this solution was 2 mM. (This value was calculated from the calibration curve of the Pt electrode shown in Fig. 3.) 2.2.5. Preparation of 500 M iodine solution in 25 mM sulphuric acid. The appropriate amount (63.5 mg) of iodine crystals were added to 500 mL of 25 mM H2SO4 in a volumetric flask applying magnetic stirring for one day. Then it was kept closed until used. 2.3. Apparatus. The experimental setup is shown in Fig. 1. The reactor was a 80 mL double walled glass beaker thermostated to 25 °C in all experiments. The usual volume of the liquid in the reactor was 50 mL.

Figure 1. Symbolic diagram of the apparatus applied in the potentiometric and in the exchange current measurements. The multimeter (MM) and the source meter (SM) were controlled by a


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personal computer (PC). In case of potentiometric experiments the source meter, the Pt reference and auxiliary electrodes played no role, and the multimeter measured the potential periodically between the Pt working and the Ag/AgCl reference, and also between the AgI and the Ag/AgCl electrodes. During exchange current measurements the polarizing current was flowing between the working and the auxiliary electrodes (Pt working and Pt auxiliary, respectively) and the overpotential  was measured between the Pt working and the Pt reference electrodes.

2.4. Electrodes. The AgI electrode was a homemade, molten-type Ag/AgI electrode, fabricated according to the method described in an earlier paper6. Pt electrodes. For calibration measurements a Pt needle electrode was used. For measuring the exchange current two more Pt electrodes were used: another Pt needle electrode as a reference to the overpotential, and a platinized Pt auxiliary electrode to set the polarizing current. One of the Pt needle electrodes was a commercial one (RADELKIS) while the other was homemade. The construction of the homemade Pt electrode was similar to that of the AgI electrode6 except that it contained a 0.5 mm diameter Pt wire. The platinized Pt electrode was taken out from a conductometer. The Ag/AgCl reference electrode was a homemade, molten-type electrode6. It was immersed in an electrolyte containing 0.1 M KCl, AgCl precipitate suspension (which protects the reference electrode from an iodide ion contamination), and 25 mM or 1 M H2SO4 (the concentration of H2SO4 was always equal to that of the solution in the reactor). The salt bridge connecting the electrolyte of the reference electrode with the reactor contained also 25 mM or 1 M H2SO4, respectively. The aim of applying the same H2SO4 concentration everywhere was to



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minimize the liquid junction potentials. The salt bridge was washed from time to time with fresh H2SO4 solutions to remove contaminations caused by previous experiments. 2.5. Measuring system. All electrode potentials were measured every 0.5 s against the Ag/AgCl reference electrode with a Keithley model 2000 multimeter (MM) equipped with a Model-2000-scanner card. Polarization measurements were carried out with a Keithley model 2410 source meter (SM). The reference electrode to the voltage generator was an unpolarized platinum wire. The current between the polarized Pt wire (working electrode) and a platinized Pt electrode (auxiliary electrode) was measured every 0.5 s with the source meter. In titration experiments, an ISMATEC REGLO DIGITAL peristaltic pump was applied to create a constant inlet of the titrant. The flow rates were calibrated by measuring the weight of the titrant removed from a small beaker during the period of titration. (The density of the 0.01 M KI solution at laboratory temperature was assumed to be 1.00 g/cm3 as a good approximation.) All devices were controlled with a computer. The double-walled reactor was thermostated to 25 °C in each experiment.

3. Measurements 3.1. Titration of hypoiodous acid with iodide. In the first experiment the potential response of Pt and AgI electrodes was recorded simultaneously while a hypoiodous acid solution was titrated with iodide. The main aim of this experiment was to test the potential response of a Pt electrode to a changing (here: decreasing) HOI concentration. The additional AgI electrode was applied to compare the response of the AgI and Pt electrodes in the course of the titration.


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The reaction vessel contained initially 50 mL of 500 M iodine solution in 25 mM H2SO4. This way the increase of [I2] by the end of the titration was relatively small (less than 10%) and the electrode potential changes were caused mostly by the order of magnitude changes in [HOI] and [I–]. 100 L of I(+1) solution in DCM was injected before the titration and this solution was titrated with a constant inlet of 10.0 mM KI solution in water (with a flow rate of 0.102 L/s). The titration was started 167 s after the injection. The initial concentration of HOI was about 45.6 M calculated from the KI consumption at the equivalence point. The titration curves are shown in Fig. 2. (The actual initial concentration was somewhat higher because the disproportionation of HOI was not taken into account in our calculation.)

Figure 2. Following the titration of HOI with I– by a) AgI and b) Pt electrodes. [HOI]0 = 45.6 M. Details of the titration are given in the Experimental part. The initial potential jump in the diagrams indicates the addition of 100 L of ( 1) solution in DCM to 500 μM iodine in 25 mM sulphuric acid. According to our digital records the inflection point of the AgI electrode potential occurred at tAgI(infl) = 2393.6 s while for the Pt electrode that happened at tPt(infl) = 2393.9 s, which values are the same within the experimental error. In addition it is interesting to observe that the two curves are practically identical except a shift along the y axis. 13 ACS Paragon Plus Environment


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3.2. Calibration of the Pt electrode for the I–/I2 and HOI/I2 redox couples. The first experiment demonstrated a qualitatively similar potential response of the AgI and Pt electrodes while HOI was titrated by I–. Next, to understand the potential response of the Pt electrode quantitatively, we measured and depicted its potential (in fact the measured electromotive force of a galvanic cell EPt,m consisting of a Pt and a Ag/AgCl reference electrode) as a function of the logarithm of [I–]

2

or

2

.

The potential of the Pt electrode was measured while adding a KI solution to an iodine solution for the I–/I2 calibration, and by adding an AgNO3 solution to an iodine solution for the HOI/I2 calibration. At low concentrations, the amount of iodide present initially in the iodine solution cannot be neglected. Therefore before starting the calibrations the amount of iodide was decreased by adding small portions of a AgNO3 solution to the iodine solution so that the potential be close to the hydrolysis limit potential on the I– or the HOI side, respectively. The initial I– and Ag+ concentration was then calculated back from the fitted calibration line so that its slope be equal to the Nernstian value of 59.16 mV/decade. The calibration curve is shown in Fig. 3. In case of the I–/I2 calibration, the reactor contained 50 mL of 500 M iodine solution in 25 mM H2SO4. Two 50 L portions of 1 mM AgNO3 were added before starting the calibration. This way the potential was close to the HLP on the I– side. The fitted value for the initial iodide concentration was [I–]0 = 0.053 M after adding the AgNO3 solution. Then a 1 mM KI solution was added in the following portions: 10, 10, 20, 40, 80, 160, 320, 640, 1280, and 2560 L. In case of the HOI/I2 calibration, the reactor contained 50 mL of 500 M iodine solution in 25 mM H2SO4. Three 50 L portions of 1 mM AgNO3 were added before starting the calibration. This way the potential was close to the HLP on the HOI side. The fitted value for the initial silver 14 ACS Paragon Plus Environment

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concentration was [Ag+]0 = 0.409 M. Then a 1 mM AgNO3 solution was added in the following portions: 10, 10, 20, 40, 80, 160, 320, 640, 1280, and 2560 L.

Figure 3. Measured EPt vs. lg(x) calibration lines of the Pt electrode. x = [I–] hydrolysis limit while x =

2

2

below the

above that limit. Extensions of the two lines cross each

other at the hydrolysis limit potential (HLP) where [I–] = [HOI]. Two peaks of the potential oscillations recorded in a BR reaction (composition is given in the caption for Fig. 10) are also shown to demonstrate that these oscillations of the Pt electrode cross the HLP.

3.3. Potential response of the Pt electrode to I2 below and above the hydrolysis limit. The Pt electrode potential was recorded while adding iodine stepwise to an iodide solution (red curve in Fig. 4) and to a HOI solution (black curve in Fig. 4). In case of the iodide solution (red curve) initially 150 L of 0.01 M KI was added to 50 mL of 25 mM H2SO4 establishing an initial iodide concentration of 29.9 M. Then 0.1 M iodine in DMSO stock solution (which contains also 2 mM iodide) was added to it in the following small, increasing portions. First 10 L was added which is not shown in the figure because there the potential jump was too high to 15 ACS Paragon Plus Environment


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fit in the figure. The iodide concentration was 30.4 M after this step. The next five portions were 10, 20, 40, 80, and 160 L. In case of the HOI solution (black curve) initially 50 mL of 31.5 M HOI solution was present in the reactor (50 L of 31.5 mM HOI solution was added to 50 mL of 25 mM H2SO4). Then the same 0.1 M iodine in DMSO solution was added to it in the same way as above (10, 10, 20, 40, 80, and 160 L).

Figure 4. Changes of the Pt potential due to iodine addition. Arrows indicate the addition of iodine to iodide (red curve) and to HOI (black curve). In each step the iodine concentration was roughly doubled. Details of the measurements are given in the Experimental part. Green curve: calculated potential drop which would be observed due the iodide contamination alone in the iodine stock.

3.4. Determining the exchange current of the Pt electrode during the titration of HOI with iodide. The titration shown in Fig. 2 was repeated while determining the exchange current in the following way. The overpotential of the Pt needle working electrode compared to the Pt reference 16 ACS Paragon Plus Environment

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electrode (see Fig. 1) was changed with an average polarizing rate of 4 mV/s. The potential difference was changed in a stepwise manner in 0.5 s steps from –10 mV to +10 mV and back (see Fig. 5a) by adjusting the current flowing between the Pt working and the platinized Pt auxiliary electrodes to establish the required potential difference. This current I was measured every 0.5 s and plotted against the overpotential  as shown in Fig. 5b. The exchange current I0 at time t was determined from the slope of the straight line fitted by a least squares method to ten I –  data points around time t, as for small  values, we have13 F 0 T

.

Figure 5. a) Polarization voltage  vs. time  diagram during the exchange current measurements. The overpotential  – the potential between the Pt working and Pt reference electrodes – was controlled by the source meter.  was held constant for 0.5 s then it jumped up or down by 2 mV. (Unfilled dots represent the starting point of the jump, while filled dots represent points after the jump.) b) The current I flowing between the Pt working and the platinized Pt auxiliary electrodes vs. the overpotential  shown at four different times t of the titration depicted in Fig. 2. The exchange current was calculated from the slope of these lines and 17 ACS Paragon Plus Environment


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depicted in Fig. 6.

The calculated exchange current I0 is then plotted against the time during the titration as can be seen in Fig. 6.

Figure 6. Exchange current (black) and Pt potential (blue) measured during the titration of HOI with iodide. Experimental conditions are the same as in Fig. 2. The method of measuring the exchange current is described in the Experimental part.

3.5. HOI versus HOIO selectivity of the Pt electrode. In the experiment shown in Fig. 7 the potential of the Pt electrode was measured while HOIO (iodous acid) was added to a HOI solution. Initially the reactor contained 50 mL of 12.6 M HOI solution in 1 M H2SO4. (The Ag/AgCl reference electrode also contained 1 M H2SO4 in this experiment.) In experiment shown in Fig. 7a) at the time indicated by the arrow 320 L of I(+3) in 96% H2SO4 was added. The

concentration right after the injection was 356.4 μM.

In experiment shown in Fig. 7b) the same volume of concentrated pure sulphuric acid but now without I(+3) (i.e. 320 L of 96% H2SO4) was added. 18 ACS Paragon Plus Environment

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Figure 7. Effect of HOIO on the Pt potential in the presence of HOI. The arrow indicates addition of 320 L of a) I(+3) dissolved in 96% H2SO4; b) 96% H2SO4 (without I(+3)).

3.6. Measuring mixed potentials in the simultaneous presence of iodine/oxyiodine species and H2O2. In this measurement, two reactors were applied (see Fig. 8): one of the reactors contained oxyiodine species and the other a H2O2 solution. In reservoir A a HOI solution was titrated with iodide with the same concentrations as shown in Fig. 2. (The initial HOI concentration was slightly different here, [HOI]0 = 48.7 M.) In reservoir B there was a 0.66 M H2O2 solution in 25 mM H2SO4. The two reservoirs were connected with a salt bridge filled with 25 mM H2SO4.Two identical Pt and AgI electrodes were used in both reservoirs and their potential was measured against the same reference electrode.



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Figure 8. Schematic drawing of the apparatus to measure mixed potentials. AgI(test), Pt(test) and AgI(pert.), Pt(pert.) indicate the wires leading from the multimeter MM to the test and the perturbing electrodes, respectively. Switches S1 and S2 could connect or disconnect the test and the perturbing electrodes.

This method of separate electrolytes and electrodes was necessary to measure mixed potentials without allowing a reaction between the iodine/oxyiodine species and H2O2. In reservoir A, HOI was titrated with iodide; this way the medium for the “test” electrode was a solution containing oxyiodine species in varying concentrations. The “perturbing” electrode was immersed into a H2O2 solution in reservoir B. Then for some intervals the Pt(test) and the Pt(pert.) or the AgI(test) and the AgI(pert.) electrodes were electrically connected and the resulting mixed potential was measured against the reference electrode. It should be realized that in this case a current is flowing through the test and the perturbing electrodes (while there is no current through the reference electrode). Thus the mixed potential – unlike the electrode 20 ACS Paragon Plus Environment

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potentials discussed previously in this paper – is not an equilibrium potential. The mixed potential is somewhere between the potentials of the test and the perturbing electrode, closer to the potential of the electrode with the larger exchange current12. The potential vs. time curves can be seen in Fig. 9.

Figure 9. Individual and mixed potentials of a) Pt and b) AgI test and perturbing electrodes. The test electrodes (black curves) were placed into reservoir A where HOI was titrated with iodide (see Fig. 2). The perturbing electrodes (red curves) were immersed into reservoir B containing a 0.66 M H2O2 solution (the same concentration we used in the BR reaction). Black and red individual potential curves are merged to a common mixed potential when the electrodes were connected.

4. Results and discussion 4.1. Following the titration of hypoiodous acid with iodide with AgI and Pt electrodes. The titrant I– flowing into the reactor reacts with HOI rapidly producing I2 to restore the iodine hydrolysis equilibrium HOI + I– + H+ ⇌ I2 + H2O. 21 ACS Paragon Plus Environment


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Another reaction, the disproportionation of HOI can also contribute to the HOI consumption. Nevertheless, regarding the main aim of this experiment, it is unimportant to know which reactions or processes are responsible for the decreasing HOI concentration which is monitored simultaneously by a AgI and a Pt electrode. The I– inflow and the disproportionation gradually diminish the HOI concentration available in the reactor and after a while an equivalence point is reached where [HOI] = [I–]. The potential vs. time diagram (Fig. 2) is the classical sigmoid type curve for both electrodes where the change of the logarithmic reactant concentrations (and consequently also the change of the measured voltage) is the fastest at the equivalence point indicated by the inflection point of the sigmoid. As can be seen the two electrodes indicate the equivalence point in the same moment within the experimental error (2393.6 s for the AgI and 2393.9 s for the Pt electrode). (It is also an important observation that the shapes of the titration curves in Fig. 2 a) and b) are the same within the experimental error. That can be explained by a nearly ideal Nernstian response of the AgI electrode which was not observed previously. The close to ideal behavior is due to the relatively high (500 micromolar) iodine concentration as will be discussed in a later publication.) It was already shown6 that a AgI electrode can be applied as an indicator electrode in such a potentiometric titration but a similar behavior of the Pt electrode is a new observation which requires further discussion. It is known that when the iodide concentration is high the Pt potential is determined by the I2/I– redox couple and the minute equilibrium HOI concentration (which can be calculated from the iodine hydrolysis and the triiodide equilibria) does not play any role. At the start of the titration, however, there is a great excess of HOI over I–, as that is indicated by the AgI electrode. Regarding the initial HOI, I2 and H+ concentrations and the iodine hydrolysis equilibrium the I– 22 ACS Paragon Plus Environment

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concentration should be below 210–10 M. It is rather improbable that such an extremely low I– concentration would be able to establish a stable electrode potential. Still, the Pt potential was quite stable even from the very beginning of the titration. 4.2. Calibration of the Pt electrode for the I–/I2 and HOI/I2 redox couples. The Pt calibration curve shown in Fig. 3 can be compared with the analogous Fig. 3 in Ref. 6 where EAgI vs. lg[I–] or lg[Ag+] calibration lines of the AgI electrode are displayed. Details of the calibration procedure are discussed in the Experimental part. I2 – I– mixtures were prepared simply by adding I– step by step to an iodine solution. The iodide and iodine concentrations were calculated considering also the tri-iodide equilibrium with KI3 = 713 M–1 [Ref. 14] . I2 + I– ⇌ I3– I2 – HOI mixtures, however, were not prepared by an analogous method (i.e. by adding HOI to an iodine solution). In theory, we could also have used the values of the HOI – I– titration curve (Fig. 2) where the HOI concentration is varying continuously in the course of the titration. In fact, such a method was already applied in a previous publication of ours6 in the calibration of the AgI electrode for [HOI]. However, here we have a new possibility, namely adding silver ions to an iodine solution which can produce HOI in the reaction: I2 + H2O + Ag+ ⇌ HOI + H+ + AgI . (That was not possible in the case of the AgI electrode which is sensitive to Ag+ ions but Pt is insensitive to Ag+.) HOI concentrations can be calculated from the amount of the added Ag+ ions regarding the AgI solubility and iodine hydrolysis equilibria. Ag+ + I– ⇌ AgI

s

Ag

–


.51 10–1 M2


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–

I2 + H2O ⇌ I– + HOI + H+

2

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5.33 10–13 M2

The decomposition of HOI was neglected because the concentration of HOI was low enough. Applying the discontinuous I2 + Ag+ method to calibrate the Pt electrode for HOI provides the following benefits compared to the continuous titration method: i) A well-defined amount of HOI is produced in each step after adding a discrete amount of Ag+ ions to the solution and we can wait until an equilibrium potential is reached. Obviously, a step-by-step (i.e. non-continuous) titration could have similar advantages but that is not suitable for the accurate determination of the equivalence point. (The equivalence point is needed to calculate the initial and other HOI concentrations of the solution.) ii) We can start with small Ag+ additions – consequently with small initial HOI concentrations – where the error due to HOI disproportionation is small. In a continuous titration of HOI by I– the initial [HOI] is large thus the HOI disproportionation is more disturbing already at the start. Moreover, around the equivalence point the potential changes too rapidly, thus potential values for low HOI concentrations become uncertain as there is not enough time to reach equilibrium. iii) By using Ag+ ions instead of I(+1) solutions in DCM, uncertainties due to a I(0) or I(+3) contamination of the I(+1) solution do not matter. The measured potentials were reproducible within 1–2 mV on the same day but the intercept values exhibited a systematic downward drift of 3–5 mV from one day to another in case of freshly prepared reference electrode and salt bridge electrolyte solutions15. This potential drift was caused by a change in the potential of the reference electrode, so to compensate that error our first idea was to shift all the measured potentials upward with the same value. The value of this correcting shift was determined by measuring an iodide calibration line with an auxiliary 24 ACS Paragon Plus Environment

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AgI electrode applying the same reference electrode. (That gave a 10.3 mV correcting term in case of Fig. 3 for example. Experimental EPt values in Fig. 3 already include that correction.) This way any effect of unknown liquid junction potentials is also eliminated. Equations for the measured logarithmic calibration lines for I–/I2 and HOI/I2 were determined so that only the intercepts were fitted to the measured points while their slope was held fixed at 59.16 mV /decade regarding the nearly ideal Nernstian response of the Pt electrode in aqueous iodine – iodide solutions8. Nernstian response also above the HLP was assumed as a working hypothesis. Applying an ordinary least squares method for fitting, the measured intercept value below the HLP was E0,Pt,m(I–/I2) HLP it was E0,Pt,m(HOI/I2)

62. ± 0.22 mV (

2

325. ± 0.36 mV ( 2

= 0.99954) and above the

= 0.99978).

Equations of the theoretical calibration lines of a Pt electrode for I–/I2 and HOI/I2 redox couples are given in the Supporting Information S1. These theoretical calibration lines are compared with the measured ones regarding their intercepts. The EMF of a galvanic cell with our Ag/AgCl reference electrode filled with 0.1 M Cl– below the HLP can be given as t(

–

2)

0,

t(

–

–

2 ) – 5 .16 mV lg

with E0,Pt,th(I–/I2) = 332.8 mV theoretical and E0,Pt,m(I–/I2)

2

325. ± 0.36 mV measured intercept

values, respectively, while above the HLP it is t(

2)

0, t (

2)

5 .16 mV  lg 2

with E0,Pt,th(HOI/I2) = 968.1 mV theoretical and E0,Pt,m(HOI/I2) = 62. ± 0.22 mV measured intercept values, respectively. As can be seen the measured I–/I2 calibration line is running 5.4 mV, and the HOI/I2 line 7.6 mV below the theoretical one. A much better agreement between the theoretical and the measured intercept values can be 25 ACS Paragon Plus Environment


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achieved, however, if the potential shift of the reference electrode is recalculated regarding the Pt potential calibration line below the HLP as a reference. The reason for this is that it is known that the Pt electrode potential is very stable and reliable in this region8. With this method all effects due to unknown liquid junction potentials are eliminated again, just as when applying an auxiliary AgI electrode. (Nevertheless, it poses the question why the correcting shift calculated with the values measured by the auxiliary AgI electrode was not as good as with the Pt. That difference is due to a small memory effect of the AgI electrode caused by some absorbed elementary iodine. That effect will be discussed in a subsequent paper.) It is also advantageous that the very same Pt electrode is applied to measure both calibration lines thus individual irregularities of the electrodes are mostly eliminated this way. (We have to remark, though, that two separate Pt electrodes usually display the same potential within a few tenths of a mV, thus “individual” effects are generally small. On the other hand, the potential difference between two AgI electrodes submerged into the same solution can easily exceed several millivolts due to the already mentioned memory effect.) Now, if both experimental lines are shifted up by an additional 7.0 mV (that additional correction is not shown in Fig. 3), then the theoretical and the corrected experimental Pt calibration lines below the HLP obviously coalesce, while above the HLP the deviation between the two lines is only 1.6 mV which is in the order of the experimental error. Thus we can conclude that for the present purpose Pt is a more reliable electrode than AgI. 4.3. Estimation of the iodine hydrolysis constant K°H based on our measurements. The equation of the theoretical calibration line is given in Supporting Information S1 in the following form: EPt(HOI/I2) / mV = 332.8 – 59.16lg

a(

o

)

+ 59.16lg

2

.

i.e. the value of E0,Pt,th(HOI/I2) depends on a(H+) and K°H. The theoretical intercept value of 26 ACS Paragon Plus Environment

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E0,Pt,th(HOI/I2) = 968.1 mV was calculated using the known hydrogen ion activity a(H+) = 0.029 M in a 25 mM sulfuric acid solution16 and a hydrolysis equilibrium constant of K°H = (5.30.1)10–13 M2 suggested by Schmitz14. This way the difference between the theoretical intercept and the corrected experimental one (the latter being 962.7 + 7.0 = 969.7 mV) is only 1.6 mV. A complete agreement between the experimental and theoretical calibration lines above the HLP can be achieved using K°H = 4.9710–13 M2 which would give a theoretical intercept of E0,Pt,th(HOI/I2)

6 . mV. Assuming a ±1 mV experimental error in the position of the

calibration line above the HLP we can estimate that K°H = (4.97 ± 0.20)10–13 M2. 4.4. Experiments showing that iodine is the reduced form in the HOI/I2 redox couple. In the previous experiments the I2 concentration was kept basically constant, and the potential of the Pt electrode was changing because of the iodide or HOI concentration changed. The next experiment demonstrates that as the potential of the Pt electrode is determined by the concentrations of two species, i.e. either by the I2/I–, or by the HOI/I2 redox couple, its potential can be modified also by changing only the iodine concentration at otherwise constant I– or at constant HOI levels (Fig. 4). Observe the opposite effect of the added iodine below and above the hydrolysis limit. When [I–] is high, increasing [I2] increases the potential in the usual way, but above the hydrolysis limit, that is where the HOI/I2 redox couple is the potential determining one, increasing [I2] decreases the potential. The green curve in Fig. 4 is a hypothetical one. It shows calculated Pt potential drops that would be caused by the iodide contamination in the iodine stock solution, disregarding the effect of iodine itself. It clearly demonstrates that the role of the iodide contamination in the observed downward steps is minor, thus the effect is due mostly to iodine.



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4.5. Measuring the exchange current of the Pt electrode. All the previous experiments (see Figs. 2–4) demonstrated that potentials above the HLP can be interpreted as HOI/I2 redox potentials. That is a reasonable interpretation as in this region the concentration of HOI is much higher than that of I–. Nevertheless, from a thermodynamic point of view potentials could be explained as a I–/I2 redox potential also above the HLP. To prove that in this region HOI is the real potential determining species we carried out exchange current experiments. The titration of HOI with iodide (shown in Fig. 2) was repeated while the exchange current was measured. The Pt potential and the simultaneously measured exchange current during the titration are shown in Fig. 6. As can be seen the latter is an asymmetric curve with a close-to-zero minimum value at the inflection point. According to the Butler-Volmer equation17,18 the exchange current I0 increases monotonically with the concentration of the oxidized and reduced species cox and cred: I0

1–

F k0 cox cred

where k0 is the mean kinetic constant of the studied redox reaction and

is the charge transfer

coefficient. So, if the Pt potential would be determined by the I–/I2 redox pair then the measured exchange current would be a monotonically increasing function of time during the whole titration process, as the iodide concentration monotonically increases during the titration (the iodine concentration is not constant but it also increases slightly). In contrast to that expectation the measured exchange current is monotonically decreasing until the equivalence point and it is increasing only when we are on the iodide side (to the right of the inflection point, for t > 2400 s). Thus we can conclude that in the region above the HLP (for t < 2400 s) the monotonic decrease


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of the exchange current supports that there the potential determining redox couple must be different. Obviously, here the oxidized species is HOI while I2 is the reduced species in the following redox equilibrium: HOI + e– + H+ ⇌ ½ I2 + H2O. On the other hand, in the region below L the behavior is “normal”: the exchange current increases with the increasing I– concentration indicating that in this region the electrode potential is determined by the I2/I– redox couple in the well known redox equilibrium: ½ I2 + e– ⇌ I –. Thus we proved that beside I2/I– also the HOI/I2 redox couple can determine the Pt electrode potential. It is logical to ask whether other oxyiodine species could also affect the potential of the Pt electrode? The next experiment intends to find an answer. 4.6. HOI versus HOIO selectivity of the Pt electrode. The previous experiments demonstrated that the HOI/I2 i.e. the I(+1)/I(0) redox pair can determine the Pt potential. It is known that in the BR reaction, however, I(+3) and I(+5) also occur and, at least in theory, they could also affect the Pt potential. No I(+5) sensitivity was reported in the literature but we have no information about the behavior of I(+3) in this respect. In an earlier experiment6 we tested whether I(+3) is able to perturb the potential of the AgI electrode and found that in the presence I(+1) I(+3) had no effect. That time we measured the Pt potential simultaneously and the result of that experiment is reported here. Initially the reactor contained 12.6 M HOI solution in 1 M H2SO4. Here the concentration of H2SO4 was much higher than in other experiments because the I(+3) stock solution was in concentrated sulphuric acid. Injection of that stock into a 25 mM H2SO4 medium would have caused a too high change in the acidity. Moreover the disproportionation of the HOIO formed



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from the injected I(+3) is slower in 1 M H2SO4. In experiment a) at the time indicated by the arrow 320 L of I(+3) stock was injected. If I(+3) has an effect on the Pt electrode then the potential should increase. Instead of the expected rise a sudden 3 mV drop was observed before a gradual increase in the potential caused by the growing HOI concentration due to the disproportionation of HOIO. Obviously the small potential drop at the injection cannot be attributed to the HOIO. The reason for this drop is most probably a change of the liquid junction potential due to the increased H2SO4 concentration. To check this assumption in experiment b) the same volume of concentrated pure sulphuric acid but now without I(+3) (i.e. 320 L of 96% H2SO4) was added. Really, the Pt potential started suddenly to decrease and finally it dropped about 6 mV. We can conclude from these experiments that HOIO has no measurable effect on the Pt potential, at least in the presence of HOI. Moreover, as the injected sample contains a considerable amount of I(+5) the experiment confirms that under these conditions the Pt electrode is not sensitive to iodate either. 4.7. The simultaneous effect of H2O2 and iodine/oxyiodine species on the potential of the Pt electrode and on the AgI electrode. In the BR reaction, iodine/oxyiodine species and H2O2 are present simultaneously. Until now it was assumed that the applied Pt and AgI electrodes respond to the iodine/oxyiodine intermediates exclusively. Here we wanted to study the perturbing effect of H2O2 on the Pt and AgI potentials. To this end we measured mixed potentials of both type (Pt or AgI) of electrodes. See the Experimental for details of the measurements. The recorded potentials are shown in Fig. 9. As can be seen, without electric connection, the test electrodes behaved as usual. However, their potential could change when they were connected with the perturbing electrode if the effect


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of H2O2 was strong enough. It can be seen that in the case of Pt electrode (Fig. 9a) the perturbing effect of H2O2 was different on the HOI and on the I– side. At the beginning of the titration when [HOI] > [I–] (above the HLP) electric connection results in a mixed potential which is clearly different from both individual potentials. The potential of the test electrode decreased by more than 100 mV. However, after the inflection point when [I–] > [HOI], the Pt potential was not perturbed significantly by H2O2. In the case of the AgI electrode (Fig. 9b) the situation was different: the mixed potential was practically equal to the potential of the AgI electrode in the reservoir with iodine species. The perturbing effect of H2O2 was negligible in this case both above and below the SLP. 4.8. Oscillations in a Briggs-Rauscher reaction followed with a AgI and a Pt electrode. Oscillations measured with the AgI and Pt electrodes are depicted in Fig. 10 a) and b), respectively. The potentials cross the SLP and the HLP, i.e. the iodide concentration periodically drops to very low levels. The potential of the AgI electrode is determined by the HOI concentration above the SLP. In Ref. 6 the maximum HOI concentration was already calculated from the AgI potential as 4–5 M. (The uncertainty is due mainly to the iodine absorbed in the AgI.) The potential of the Pt electrode above the HLP is a mixed potential, however, which – beside the HOI/I2 concentrations – is strongly affected also by the H2O2 concentration. Increasing HOI concentration would increase the potential while increasing H2O2 concentration would decrease the potential. (That is without the disturbing effect of H2O2 the potential peaks of the Pt electrode were even higher.) It is remarkable that the Pt potential could rise about 50 mV above the HLP in spite of the presence of H2O2. It seems that when H2O2 and iodine species are in the



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same reservoir in this environment the H2O2 has less effect on the Pt potential. That might be explained by the fact that iodine and iodide are able to adsorb on the surface of the Pt electrode19,20.

Figure 10. Oscillations followed simultaneously by a) AgI and b) Pt electrodes in the BriggsRauscher reaction. Initial conditions (after adding all components): [H2SO4] = 25 mM, [KIO3] = 40 mM, [H2O2] = 0.66 M, [MA] = 50 mM, [Mn2+] = 6.5 mM.

5. Conclusions Our experiments demonstrate that in a mixture of oxyiodine species above the HLP the potential of a Pt electrode is controlled not by the I2/I– but by the HOI/I2 redox couple exclusively as that is shown by our exchange current measurements. Based on the intercept of the Pt calibration line we have an estimated value of the iodine hydrolysis constant as K°H = (4.97 ± 0.20)10–13 M2 in an acceptable agreement with the value of K°H

(5.3 ± 0.1)10–13 M2 suggested by Schmitz14. To check the validity of this result the

statistical error could be decreased even further with more careful measurements (e.g. by 32 ACS Paragon Plus Environment

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applying a Pt(I–/I2) reference electrode with a saturated iodine solution or by increasing the number of measured points) but before doing that it would be important to estimate the systematic errors of this new KH determination method (e.g. the errors introduced by our working hypothesis assuming a theoretical Nernstian response of the Pt electrode). Our mixed potential measurements showed that in a system containing oxyiodine species and H2O2 simultaneously, a Pt potential above the HLP clearly indicates the presence of HOI but its concentration cannot be calculated from the electrode potential alone as that is also affected by [H2O2] as well. This prevents calculating component concentrations in the BR reaction oscillations whenever the Pt potential is above the HLP. On the other hand, H2O2 does not perturb the potential of a AgI electrode significantly.

Corresponding Author Maria Wittmann, email: [email protected]

Acknowledgement. This work was partially supported by OTKA grant K104666.

Supporting Information Available: S1. Pt electrode: theoretical calibration lines below and above the HLP S2. AgI electrode: calculated surface and bulk I– concentrations above the SLP



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Gottardi, W. Redox-potentiometric/titrimetric analysis of aqueous iodine solutions.

Fresenius J. Anal. Chem. 1998, 362, 263-269. (9)

Handbook of reference electrodes; Inzelt, G.; Lewenstam, A.; Scholz, F. Eds. Springer, 34 ACS Paragon Plus Environment

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2013, p. 121. (10) Katsounaros, I.; Schneider,W. B.; Meier, J.C.; Benedikt, U. P.; Biedermann, U.; Auer A. A.; Mayrhofer, K. J. J. Hydrogen peroxide electrochemistry on platinum: towards understanding the oxygen reduction reaction mechanism. Phys. Chem. Chem. Phys. 2012, 14, 7384–7391. (11) Electrochemical Dictionary; Bard, A. J.; Inzelt, Gy.; Scholz, F. Eds. Springer, 2012, p. 740. (12) Park, J. H.; Zhou, H.; Percival, S. J.; Zhang, B.; Fan, F-R. F.; Bard, A. J. Open circuit (mixed) potential changes upon contact between different inert electrodes – size and kinetic effects. Anal. Chem. 2013, 85, 964–970. (13) Bockris, J. ’M.; eddy, A. . N. Modern electrochemistry: an introduction to an interdisciplinary area. Plenum Press, New York, 1970, Vol. 2, Chap. 8. (14) Electronic supplement of Schmitz, G. Inorganic reactions of I(+1) in acidic solutions. Int. J. Chem. Kinet. 2004, 36, 480-493. (15) The observed potential drift is probably due to a combination of the decreasing chloride ion and the increasing iodine concentrations in the electrolyte of the reference electrode. These concentration changes are due a slow diffusional transport via the salt bridge. Liquid junction potentials can play a less important role here because the sulfuric acid concentration is 25 mM everywhere and the ionic mobilities of the K+ and Cl– ions are roughly the same. As the drift was more pronounced right after filling fresh solutions into the salt bridge and to the reference electrode, we waited one-two days before making sensitive measurements when the rate of the drift became smaller. The error caused by the overall potential shift was taken into account as it is described in the main text. (16) Clegg, S. L.; Rard, J. A.; Pitzer, K. S. Thermodynamic properties of 0–6 mol kg–1 aqueous 35 ACS Paragon Plus Environment


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sulfuric acid from 273.15 to 328.15 K. J. Chem. Soc. Faraday Trans. 1994, 90, 1875-1894. (17) Inzelt, G. Butler-Volmer equation. In Electrochemical Dictionary; Bard, A. J., Inzelt, G., Scholz, F. Eds.; 2nd Edition, Springer, Heidelberg, 2012, pp 90-91. (18) Lecture 13: Butler-Volmer equation. Notes by ChangHoon Lim. http://ocw.mit.edu/courses/chemical-engineering/10-626-electrochemical -energy-systemsspring-2014/lecture-notes/MIT10_626S14_S11lec13.pdf (accessed Aug. 30, 2016) (19) Breiter, M. W. Voltammetric study of halide ion adsorption on platinum in perchloric acid solutions. Electrochim. Acta 1963, 8, 925–935. (20) Podlovchenko, B. I.; Kolyadko, E. A. Variations in the charge and open-circuit potential of a platinum electrode during adsorption of iodine and iodide ions. Russian J. Electrochem. 2000, 36, 1268–1274.


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Figure 1. Symbolic diagram of the apparatus applied in the potentiometric and in the exchange current measurements. The multimeter (MM) and the source meter (SM) were controlled by a personal computer (PC). In case of potentiometric experiments the source meter, the Pt reference and auxiliary electrodes played no role, and the multimeter measured the potential periodically between the Pt working and the Ag/AgCl reference, and also between the AgI and the Ag/AgCl electrodes. During exchange current measurements the polarizing current was flowing between the working and the auxiliary electrodes (Pt working and Pt auxiliary, respectively) and the overpotential η was measured between the Pt working and the Pt reference electrodes. 82x84mm (300 x 300 DPI)

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Figure 2. Following the titration of HOI with I– by a) AgI and b) Pt electrodes. [HOI]0 = 45.6 µM. Details of the titration are given in the Experimental part. The initial potential jump in the diagrams indicates the addition of 100 µL of I(+1) solution in DCM to 500 µM iodine in 25 mM sulphuric acid. According to our digital records the inflection point of the AgI electrode potential occurred at tAgI(infl) = 2393.6 s while for the Pt electrode that happened at tPt(infl) = 2393.9 s, which values are the same within the experimental error. 82x59mm (300 x 300 DPI)



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Figure 2. Following the titration of HOI with I– by a) AgI and b) Pt electrodes. [HOI]0 = 45.6 µM. Details of the titration are given in the Experimental part. The initial potential jump in the diagrams indicates the addition of 100 µL of I(+1) solution in DCM to 500 µM iodine in 25 mM sulphuric acid. According to our digital records the inflection point of the AgI electrode potential occurred at tAgI(infl) = 2393.6 s while for the Pt electrode that happened at tPt(infl) = 2393.9 s, which values are the same within the experimental error. 82x59mm (300 x 300 DPI)


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Figure 3. Measured EPt vs. lg(x) calibration lines of the Pt electrode. x = [I–] / √[I2] below the hydrolysis limit while x = [HOI] / √[I2] above that limit. Extensions of the two lines cross each other at the hydrolysis limit potential (HLP) where [I–] = [HOI]. Two peaks of the potential oscillations recorded in a BR reaction (composition is given in the caption for Fig. 10) are also shown to demonstrate that these oscillations of the Pt electrode cross the HLP. 82x61mm (300 x 300 DPI)



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Figure 4. Changes of the Pt potential due to iodine addition. Arrows indicate the addition of iodine to iodide (red curve) and to HOI (black curve). In each step the iodine concentration was roughly doubled. Details of the measurements are given in the Experimental part. Green curve: calculated potential drop which would be observed due the iodide contamination alone in the iodine stock. 82x59mm (300 x 300 DPI)


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Figure 5. a) Polarization voltage η vs. time τ diagram during the exchange current measurements. The overpotential η – the potential between the Pt working and Pt reference electrodes – was controlled by the source meter. η was held constant for 0.5 s then it jumped up or down by 2 mV. (Unfilled dots represent the starting point of the jump, while filled dots represent points after the jump.) b) The current I flowing between the Pt working and the platinized Pt auxiliary electrodes vs. the overpotential η shown at four different times t of the titration depicted in Fig. 2. The exchange current was calculated from the slope of these lines and depicted in Fig. 6. 82x66mm (300 x 300 DPI)



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Figure 5. a) Polarization voltage η vs. time τ diagram during the exchange current measurements. The overpotential η – the potential between the Pt working and Pt reference electrodes – was controlled by the source meter. η was held constant for 0.5 s then it jumped up or down by 2 mV. (Unfilled dots represent the starting point of the jump, while filled dots represent points after the jump.) b) The current I flowing between the Pt working and the platinized Pt auxiliary electrodes vs. the overpotential η shown at four different times t of the titration depicted in Fig. 2. The exchange current was calculated from the slope of these lines and depicted in Fig. 6. 82x67mm (300 x 300 DPI)


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Figure 6. Exchange current (black) and Pt potential (blue) measured during the titration of HOI with iodide. Experimental conditions are the same as in Fig. 2. The method of measuring the exchange current is described in the Experimental part. 82x61mm (300 x 300 DPI)



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Figure 7. Effect of HOIO on the Pt potential in the presence of HOI. The arrow indicates addition of 320 µL of a) I(+3) dissolved in 96% H2SO4; b) 96% H2SO4 (without I(+3)). 82x59mm (300 x 300 DPI)


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Figure 7. Effect of HOIO on the Pt potential in the presence of HOI. The arrow indicates addition of 320 µL of a) I(+3) dissolved in 96% H2SO4; b) 96% H2SO4 (without I(+3)). 82x59mm (300 x 300 DPI)



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Figure 8. Schematic drawing of the apparatus to measure mixed potentials. AgI(test), Pt(test) and AgI(pert.), Pt(pert.) indicate the wires leading from the multimeter MM to the test and the perturbing electrodes, respectively. Switches S1 and S2 could connect or disconnect the test and the perturbing electrodes. 82x87mm (300 x 300 DPI)


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Figure 9. Individual and mixed potentials of a) Pt and b) AgI test and perturbing electrodes. The test electrodes (black curves) were placed into reservoir A where HOI was titrated with iodide (see Fig. 2). The perturbing electrodes (red curves) were immersed into reservoir B containing a 0.66 M H2O2 solution (the same concentration we used in the BR reaction). Black and red individual potential curves are merged to a common mixed potential when the electrodes were connected. 82x59mm (300 x 300 DPI)



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Figure 9. Individual and mixed potentials of a) Pt and b) AgI test and perturbing electrodes. The test electrodes (black curves) were placed into reservoir A where HOI was titrated with iodide (see Fig. 2). The perturbing electrodes (red curves) were immersed into reservoir B containing a 0.66 M H2O2 solution (the same concentration we used in the BR reaction). Black and red individual potential curves are merged to a common mixed potential when the electrodes were connected. 82x59mm (300 x 300 DPI)


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Figure 10. Oscillations followed simultaneously by a) AgI and b) Pt electrodes in the Briggs-Rauscher reaction. Initial conditions (after adding all components): [H2SO4] = 25 mM, [KIO3] = 40 mM, [H2O2] = 0.66 M, [MA] = 50 mM, [Mn2+] = 6.5 mM. 82x63mm (300 x 300 DPI)



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Figure 10. Oscillations followed simultaneously by a) AgI and b) Pt electrodes in the Briggs-Rauscher reaction. Initial conditions (after adding all components): [H2SO4] = 25 mM, [KIO3] = 40 mM, [H2O2] = 0.66 M, [MA] = 50 mM, [Mn2+] = 6.5 mM. 82x65mm (300 x 300 DPI)


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