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Impact of Various Acids and Bases on the Voltammetric Response of Platinum Group Metal Oxides.

1. Introduction

Platinum metal oxides such as ruthenium dioxide (Ru[O.sub.2]) and iridium dioxide (Ir[O.sub.2]) are known for their electrical resistivity at room temperature and stability in aggressive solutions containing hydrofluoric acid [1]. Their redox chemistry makes them attractive for applications such as pH sensors [2-8] and supercapacitors [4, 9].

Platinum metal oxides, in contact with aqueous solutions, are able to adsorption water, whereby a carpet of hydroxyl (OH) groups is formed by proton displacement during anodic charging. Two sorts of hydrogen groups exist: OH with bridging oxygen atoms which acts as Lewis acid or base, and OH that works as a Bronsted base [10]. Protons from the solution can be exchanged with the electrode surface by following simplified equation:

[Ru.sup.IV] [O.sub.2] + [e.sup.-] + [H.sup.+] [??] [Ru.sup.III] O (OH) (1)

Using Nernst's equation, the redox potential of the Ru[O.sub.2] electrode depends on the pH value. By definition in the solid state, the activities of Ru(III) and Ru(IV) approach a = 1, so that (3) is received.

[mathematical expression not reproducible]. (2)

At 25[degrees]C,

E = [E.sup.0] - 0.059 x pH. (3)

The thermodynamically calculated standard potential equals [E.sup.0] = 0.94 V [11]. According to the above equations, Ru[O.sub.2] electrodes can be utilized for a potentiometric sensor.

Usually, buffer solutions are used to investigate the pH sensitivity, whereas the impact of ionic strength has scarcely been considered in the literature so far. At present, the redox behavior of Ru[O.sub.2] in aqueous solutions is not fully understood. We present new findings on the impact of various acids and bases on the performance of Ru[O.sub.2] electrodes by the help of cyclic voltammetry (CV), open circuit potential (OCP) measurements, and electrochemical impedance spectroscopy. As well, the hydrogen insertion reaction into the lattice structure of Ru[O.sub.2] electrodes is considered.

2. Instruments and Methods

Electrodes were prepared by thermal spray pyrolysis from commercially available ruthenium(III) chloride hydrate (Sigma-Aldrich). The substrate, titanium foil (Ti) of 0.05 mm thickness (Ankuro Int. GmbH), was pretreated with abrasive paper and degreased with acetone to ensure good adhesion of the oxide layer. The Ru[Cl.sub.3] x x[H.sub.2]O was solved in acetone and then decomposed on the substrate in a furnace in air at 500[degrees]C for 2 h. Our previous investigations by thermogravimetric analysis revealed that the transition temperature of oxide formation is at about 360[degrees]C. The active area of the sensor was 1 [cm.sup.2], and the thickness of the oxide layer was about 2 [micro]m.

The electrochemical measurements were carried out using a three-electrode arrangement, a potentiostat/galvanostat (EC301 Stanford Research System, Inc.), and a frequency response analyzer (Solartron SI 1250). The reference electrode (RE) was a reversible hydrogen electrode (RHE), the counterelectrode (CE) was a platinum-foil, and the working electrode (WE) was the above-mentioned Ru[O.sub.2] electrode. As electrolytes, 70 mL of 1-molar solutions of sulfuric acid ([H.sub.2]S[O.sub.4]), hydrochloric acid (HCl), phosphoric acid ([H.sub.3]P[O.sub.4]), nitric acid (HN[O.sub.3]), sodium hydroxide (NaOH), and potassium hydroxide (KOH) were used. The temperature was 25 [+ or -] 0.5[degrees]C.

3. Results and Discussion

3.1. Open Circuit Potential in Different Solutions. The known problems of Ru[O.sub.2] electrodes in aqueous solutions are (i) potential drift and (ii) hysteresis. With respect to the pH sensitivity, the OCP of a Ru[O.sub.2] electrode were observed in 1molar acids andbases for a period of several days (see Table 1).

The titanium support plays no role, as far as Ru[O.sub.2] is coated sufficiently thick. The quality of the electrode was proved by the help of microscopy and impedance spectroscopy. Any weakly conducting layer of Ti[O.sub.2] on the substrate material, which might be formed by corrosion under an insufficient Ru[O.sub.2] coating, causes a dramatic increase of the resistance of the electrode within several days. With respect to quality assurance, any undesired peeling off of the Ru[O.sub.2] coating can easily be checked under a magnifying glass.

In [H.sub.2]S[O.sub.4], the Ru[O.sub.2] electrode provided a stable potential of about 920 mV RHE, which is in good agreement with the thermodynamically calculated standard potential (940 mV [11], (1)). The theoretical pH values in Table 1 were calculated using activity coefficients [gamma] from the literature [12], pH = -log([gamma]c), for c = 1 mol/L.

Unfortunately, the potentials in HCl, HN[O.sub.3], and alkali hydroxides vary dramatically depending on the ionic composition of the solution. According to the p[K.sub.a] values, HCl behaves as an extremely strong acid, followed by [H.sub.2]S[O.sub.4] and HN[O.sub.3], whereas [H.sub.3]P[O.sub.4] appears moderately strong. In contrast to the expectation that the strongest acid should have the highest potential, Table 1 reveals no clear trend, because the solution is determined not only by the protons, but also by the total ionic strength of all ions present. For example, Ru[O.sub.2] shows a significant sensitivity to chloride. Therefore, any empirical correlation between acidity (p[K.sub.a] value) and the OCP remains vague, as shown in Figure 1.

Chloride ions play a special role, as they interact with the Ru[O.sub.2] lattice, so that they change the open circuit potential as well as the slope dE/dpH and the cyclic voltammogram (Figure 3). Since chloride and chlorine compounds are found in municipal wastewater, the conventional potentiometric pH measurements using metal oxides are difficult.

3.2. Cyclic Voltammetry in Acidic Solutions. With respect to clarifying the electrochemical behavior of Ru[O.sub.2] in different acids, cyclic voltammetry was applied. The investigation was to show changes in the voltammograms depending on the pH value. Figure 2 presents the fundamental difference between the cyclic voltammogram of a Ru[O.sub.2] electrode in [H.sub.2]S[O.sub.4] in contrast to NaOH.

In [H.sub.2]S[O.sub.4] (pH 0.1), 14 different peaks can be distinguished besides the hydrogen and oxygen evolution reaction in the potential range between 0 and 1.4 V RHE.

At E [approximately equal to] 0 V RHE the hydrogen adsorption reaction occurs on the onefold coordinatively unsaturated ruthenium site (in short: 1f-cus Ru) and the bridging oxygen atom. Above 0.4 V RHE the oxidation of the surface takes place. In acidic solutions the redox transitions of Ru(III)-Ru(IV)-Ru(V) [right arrow] ... [right arrow] Ru(VIII) appear to be the predominant oxidation states:

2Ru[O.sub.2] + 2[H.sup.+] 2[e.sup.-] [??] [Ru.sub.2][O.sub.3] + [H.sub.2]O (4)

[mathematical expression not reproducible] (5)

[mathematical expression not reproducible] (6)

Ru[O.sub.4.sup.-] + [e.sup.-] [??] Ru[O.sub.4] (7)

Interestingly, different acids produce clearly different voltammograms (Figure 3).

Oxygen Reduction. The shapes of the anodic cycles are quite similar in different acids. However, the cathodic ramp is substantially altered by the presence of chloride in HCl. In HCl, the oxygen evolution reaction occurs at a lower potential than in the other acids. The oxygen evolution at first appears in HCl and then shifts to HN[O.sub.3], [H.sub.2]S[O.sub.4], and finally [H.sub.3]P[O.sub.4], whereby the electrolyte resistance plays a minor role. These low overpotentials are known for the chlorine evolution reaction in the chloralkali electrolysis with dimensionally stable anodes (Ti[O.sub.2]/Ru[O.sub.2], DSA). According to the Deacon process, traces of chlorine might be formed by reaction with anodically generated oxygen.

4HCl + [O.sub.2] [right arrow] 2[Cl.sub.2] + 2[H.sub.2]O (8)

The oxygen reduction peak at about 1.1V RHE was attributed to Ru(VIII), especially to the species Ru[O.sub.4] in the measured solution, based on in situ reflectance spectroscopy in 0.5 M [H.sub.2]S[O.sub.4] by Kotz et al. [13]. However, Ru[O.sub.4] is formed at high positive potentials during oxygen evolution and appears to be the main corrosion product in acidic solutions (see (7)). The formation of Ru[O.sub.4] in HCl is attributed to the shifted potential and therefore earlier oxygen evolution reaction.

Hydrogen Adsorption. By extending the potential window to more negative values, the hydrogen adsorption region appears more clearly and the peak heights of the anodic hydrogen oxidation increase. Starting the voltammogram at -0.1 V RHE, the largest current flows in [H.sub.2]S[O.sub.4]. Surprisingly, [H.sub.3]P[O.sub.4] follows, comparable to HCl, whereas HN[O.sub.3] reveals the weakest current (due to the smallest hydrogen evolution current).

We conclude that hydrogen oxidation current depends on the extent of the hydrogen evolution, which can be controlled by the potential window. Assuming that the potential window of 1.53 V leads to distinct peaks in 1M [H.sub.2]S[O.sub.4], the required potential difference was not reached in 1 M HN[O.sub.3] (about 1.48 V). About 1.4 V in 1M HCl is enough to show a pronounced voltammogram of ruthenium oxide.

The oxygen reduction reaction seems to be unaffected by the cathodically altered potential window.

3.3. Cyclic Voltammetry in Alkaline Solutions. In alkaline solution, the predominant oxidation states read Ru(III) [right arrow] Ru(IV) [right arrow] Ru(VI) [right arrow] Ru(VII). At high potentials in the region of oxygen evolution, Ru(VIII) is generated and can be reduced in the solution at lower potentials. The redox reactions can be expressed as

Ru[O.sub.4.sup.2-] + 4[H.sup.+] + 2[e.sup.-] [??] Ru[O.sub.2] + 2[H.sub.2]O (9)

Ru[O.sub.4.sup.-] + [e.sup.-] [??] Ru[O.sub.4.sup.2-] (10)

The cyclic voltammograms of Ru[O.sub.2] in NaOH and KOH are shown in Figure 4.

Using a potential window of 0 to 1.5 V, the hydrogen oxidation region at E [approximately equal to] 0 V is not pronounced. The shape of the voltammograms remains the same. However, there are differences in the peak heights, that is, for the Ru(VI)/Ru(VII) couple at 1.3 V. If the potential window is extended to more negative potentials, the hydrogen oxidation region becomes more intense, and peak height of the Ru(III) [right arrow] Ru(IV) oxidation at 0.6 V changes too. Note that in acids, this peak was not changed (or only in a negligible extent).

3.4. Impedance Spectroscopy in Acidic and Alkaline Solutions. The Ru[O.sub.2] electrode was characterized in different acidic and alkaline solutions by impedance spectroscopy in a frequency range from 65 kHz to 10 Hz. The electrochemical activity of the oxide-electrolyte interface is evaluated by the help of capacitance [14].

[mathematical expression not reproducible]. (11)

Figure 5 shows that the Ru[O.sub.2] electrode exhibits the largest interface activity in [H.sub.2]S[O.sub.4] (1450 [micro]F/[cm.sup.2] at 10 Hz), whereas the capacitance in [H.sub.3]P[O.sub.4] is the smallest (650 [micro]F/[cm.sup.2] at 10 Hz). HCl, HN[O.sub.3], and NaOH range is around 1250 [micro]F/[cm.sup.2].

Interface capacitance allows distinguishing strong and weak acids. Although the p[K.sub.a] of HCl is the highest, the capacitance is less pronounced than that of [H.sub.2]S[O.sub.4]. As seen with the cyclic voltammograms, the chloride ions compete with hydroxide sites in the Ru[O.sub.2] lattice. Other ions present in the acids do not affect the behavior of the electrode, that is, sulfate, phosphate, and nitrate.

The Nernst slope, which generally deviates from 59 mV [pH.sup.-1] (25[degrees]C), is nearly independent of dissolved anions in the solution (such as S[O.sub.4.sup.2-], P[O.sub.4.sup.2-], and N[O.sub.3.sup.-]). However, commercial Ru[O.sub.2] resistive pastes, which contain PbO, exhibit a slope which depends on different anions significantly [15].

3.5. Proton Insertion in Ruthenium Dioxide. As shown in Figure 3, the extension of the cathodic potential window leads to a significant increase in hydrogen oxidation current at E [approximately equal to] 0 V. This is consistent with the fact that Ru[O.sub.2] forms a quasi-hydrogen electrode, once a negative potential has been applied. The formally adsorbed hydrogen (even in the form of hydroxide species) is anodically oxidized according to the reversed equation (1). Cyclic voltammograms in the same solution with different starting potentials are shown in Figure 6.

A more cathodic electrode potential (in the hydrogen evolution direction) leads to (i) larger hydrogen oxidation currents and (ii) a shift of hydrogen oxidation towards more positive potentials. The impact of hydrogen on the behavior of a ruthenium oxide electrode was already mentioned in 1974 by Galizzioli et al. [11]. Continuous hydrogen evolution leads to the collapse of the atomic layers under the surface, but there is no lasting effect.

Michell et al. reported that hydrogen can adsorb in the lattice of the ruthenium oxide. This protonation of the oxide lattice leads to the formation of a large porous anodic film, which might break Ru-O bonds [16-18]. Today, we imagine proton insertion in Ru[O.sub.2] by the dissociative adsorption of water [9] and bridged OH groups [10].

(1) In acid solution, the electrode potential increases with rising pH, in the same way that is known of the hydrogen oxidation at a hydrogen electrode. Protons are released by the dissociative adsorption of water and superacid OH groups. Simply, by the help of rutile lattice sites [Ru], the potential determining surface process at the more negatively charged Ru[O.sub.2] electrode reads

[mathematical expression not reproducible] (12)

Formally: [Ru] [H.sub.2] [??] [Ru] + 2[H.sup.+] + 2[e.sup.-]

It is unclear whether Ru(II) exists during the anodic oxidation of hydrogen near 0 V according to (12). Figure 6 shows that the larger the potential is displaced in the negative direction, the higher the reversible oxidation and reduction currents of the Ru(III)/Ru(IV) couple at 0.5 V are.

(2) In alkaline solution, the electrode potential decreases with rising pH, in the same way that is known of the oxygen reduction at an oxygen electrode. By the dissociative adsorption of water, hydroxide sites are formed and bound in ruthenium cluster ions.

[mathematical expression not reproducible] (13)

Formally: [Ru] [O.sub.2] + 4[e.sup.-] + 2[H.sub.2]O [??] [Ru] + 4O[H.sup.-]

If the potential is driven far in the positive direction, the oxygen reduction at 1.4 V and the Ru(IV) [right arrow] Ru(III) reduction at 0.4 V get stronger. On the other hand, further oxidation states such as Ru(IV)/Ru(V) seem to be unaffected by the hydrogen insertion.

3.6. Diffusion Control. The fact that the hydrogen adsorption is a diffusion controlled reaction can be seen in Figure 7. The hydrogen region is clearly affected by stirring the solution, because the exchange of hydrogen ions between Ru[O.sub.2] surface and solution is disturbed. This affects to a certain extent the oxidation of the electrode at 0.8 V, because [H.sup.+] is involved in the Ru(III)/Ru(IV) redox reaction.

4. Conclusion

This work compiles the impact of different acids and hydroxide solutions on the ruthenium dioxide electrode. With respect to pH measurements, the determination of the OCP appears to be problematic in many respects, because the electrode potential varies in time and can hardly be reproduced due to hysteresis and aging effects.

Depending on the ionic strength, solutions even having the same pH and the same concentration cause different electrode potentials. Chloride ions affect strongly the behavior of the electrode. The cyclic voltammograms in acid solutions reveal alterations of the oxidation and reduction peaks due to interfering ions. In alkaline solution, the electrode seems to be affected by the potential window only.

We also demonstrated the hydrogen insertion into the bulk of the Ru[O.sub.2], forming [Ru]-OH, which plays a role in the pH-dependent Ru(III)/Ru(IV) redox couple, whereas higher oxidation states remain unaffected.

This study shows that further work has to be done to understand in detail the impact of interfering ions on the pH response of a Ru[O.sub.2] electrode.

Conflicts of Interest

The authors declare that they have no conflicts of interest.


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Sebastian Chalupczok (iD), Peter Kurzweil (iD), and Helmut Hartmann

Electrochemistry Laboratory, University of Applied Sciences, Kaiser-Wilhelm-Ring 23, 92224 Amberg, Germany

Correspondence should be addressed to Peter Kurzweil;

Received 30 August 2017; Accepted 9 November 2017; Published 23 January 2018

Academic Editor: Davood Nematollahi

Caption: Figure 1: Fuzzy correlation of acid strength p[K.sub.a] and open circuit potential of a Ru[O.sub.2] electrode in different acids and bases.

Caption: Figure 2: Cyclic voltammogram of a Ti/Ru[O.sub.2] electrode in 1M [H.sub.2]S[O.sub.4] and 1M NaOH at 25[degrees]C. Counterelectrode: Pt sheet. Scan rate: 100mV/s.

Caption: Figure 3: (a) Cyclic voltammograms of acids at 25[degrees]C. Impact of an extended potential window: solid lines: 0 to 1.5 V; dashed lines: -0.1 to 1.5 V RHE. Scan rate: 100 mV/s. (b) Full scale to show the oxygen reduction (Ru[O.sub.4]) and the impact of chloride at 1.1 V RHE.

Caption: Figure 4: Cyclic voltammograms of alkali hydroxide solutions (red: 1-molar KOH; black: 1-molar NaOH). Impact of an extended potential window (solid lines: 0 to 1.5 V; dashed lines: -0.1 and -0.2 to 1.5 V). Scan rate: 100 mV/s.

Caption: Figure 5: Frequency response of capacitance of a Ti/Ru[O.sub.2] electrode (versus RHE) in different solutions at 25[degrees] C.

Caption: Figure 6: Increasing oxidation of adsorbed hydrogen (by hydrogen insertion into bulk Ru[O.sub.2]) in a more and more cathodic potential window (hydrogen evolution) and more positive potentials (oxygen evolution). Picture inside: Ru(IV) [right arrow] Ru(III) reduction at 0.4 V.

Caption: Figure 7: Hydrogen adsorption at a Ti/Ru[O.sub.2] electrode in a stirred (300 [min.sup.-1]) and unstirred pH 7 buffer solution. CE: platinum sheet.
Table 1: Drift of the OCP (E versus RHE) of a Ru[O.sub.2] electrode
in various acids and bases at 25[degrees]C in the course of several
days. For comparison: negative decadic logarithm of the dissociation
constants of acids (p[K.sub.a]) and bases (p[K.sub.b]) from [12.].
Theoretical potential of the reversible hydrogen electrode:
V RHE = 0.059 pH (25[degrees]C).

Solution                       pH              pK         Day 1
                          (calculated)   (25[degrees]C)   (E/mV)

1 M [H.sub.2]S[O.sub.4]       0.11             -3          920
1 M HCl                       0.13             -7          815
1 M [H.sub.3]P[O.sub.4]       1.08           +2.13         866
1 M HN[O.sub.3]               0.13           -1.37         844
1 M NaOH                      13.8            -1.1         702
1 M KOH                       13.8           -0.56         846

Solution                  Day 2    Day 3
                          (E/mV)   (E/mV)

1 M [H.sub.2]S[O.sub.4]    923      920
1 M HCl                    697      850
1 M [H.sub.3]P[O.sub.4]    889      864
1 M HN[O.sub.3]            806      942
1 M NaOH                   710      912
1 M KOH                    972      902
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Title Annotation:Research Article
Author:Chalupczok, Sebastian; Kurzweil, Peter; Hartmann, Helmut
Publication:International Journal of Electrochemistry
Date:Jan 1, 2018
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