Mechanism of tetrachloroplatinate( II ) oxidation by hydrogen peroxide in hydrochloric acid solution †

Oxidation of tetrachloroplatinate( II ) by hydrogen peroxide in hydrochloric acid was studied by UV-Vis spectrophotometry. Oxidation takes place via two parallel reactions with hypochlorous acid and hydrogen peroxide, respectively, according to the overall rate law d[Pt( IV )]/d t = ( k 0 + k H 2 O 2 [Pt( II )])[H 2 O 2 ]. For oxidation of [PtCl 4 ] 2 − at relatively low concentrations, [PtCl 4 ] 2 − ≪ 0.5 mM, hypochlorous acid formation is fast relative to the oxidation of [PtCl 4 ] 2 − by hydrogen peroxide, as a result of the rate determining reaction H 2 O 2 + H + + Cl − → HOCl + H 2 O, resulting in a rate law d[Pt( IV )]/d t = k 0 [H 2 O 2 ] with a value k 0 = (8 ± 2) × 10 − 7 s − 1 at 35 °C. For concentrations of [PtCl 4 ] 2 − > 0.5 mM, oxidation by hydrogen peroxide becomes dominant, resulting in the pseudo-ﬁ rst order rate law d[Pt( IV )]/d t = k H 2 O 2 [Pt( II )][H 2 O 2 ] with the value k H 2 O 2 = (1.5 ± 0.1) × 10 − 2 M − 1 s − 1 at 35 °C. The ﬁ nal oxidation product is a mixture of [PtCl 5 (H 2 O)] − and [PtCl 6 ] 2 − , with [PtCl 6 ] 2 − formed as a result of [PtCl 4 ] 2 − assisted chloride anation reactions.


Introduction
Oxidation of Pt(II) square-planar complexes by hydrogen peroxide has been exploited in many areas of research, particularly as a strategy towards the design of new complexes. 1Many of these studies rely on the formation of hydroxido complexes by the oxidation with hydrogen peroxide, which provides more stability and control.The square-planar configuration of the original Pt(II) complex is retained furnishing a Pt(IV) product with new ligands coordinated trans to each other.For instance, oxidation of [PtCl 4 ] 2− by hydrogen peroxide in water yields trans-[PtCl 4 (OH) 2 ] 2− quantitatively according to eqn (1). 2 195 Pt NMR indicates that the trans coordinated hydroxido ligands originate from hydrogen peroxide and solvent water respectively. 3Inert hydroxido ligands can be protonated after oxidation to render the aqua ligands that are labile, promoting substitution reactions. 4y comparison, oxidation of [PtCl 4 ] 2− by hydrogen peroxide in acidic medium yields trans-[PtCl 4 (H 2 O) 2 ], while relatively fast Pt(II) assisted ligand scrambling reactions cause a redistribution of the oxidation product(s) to form a mixture of [PtCl 6−n (H 2 O) n ] 2−n (n = 0-4) complexes. 5The large-scale separation of platinum from other platinum group metals (PGMs) depends, amongst other factors, on the efficient oxidation of Pt(II) to Pt(IV) in solution.The oxidation states of the various PGMs dissolved in the hydrochloric acid process solutions are manipulated to allow for their separation by inter alia solvent extraction (SX), oxidative distillation and/or classical ionexchange methods. 5Although hydrogen peroxide is not used in the refining industry, as part of ongoing work in this context, we examined in detail the oxidation of [PtCl 4 ] 2− by hydrogen peroxide in hydrochloric acid as a benchmark system.In the oxidation of Pt(II) to Pt(IV) in solution, Pearson and Basolo proposed a reaction mechanism involving Pt(II) assisted ligand exchange more than 50 years ago. 6This mechanism was later revised to account for direct formation of [PtCl 6 ] 2− from trans-[PtCl 4 (H 2 O) 2 ] in a chloride rich solution. 4he origin of the difference in behaviour between trans-Pt(IV) aqua and hydroxido analogues is attributed to the labile trans aqua ligands which are susceptible to bond breakage. 7Kinetic studies dealing with oxidation of Pt(II) complexes by hydrogen peroxide have been neglected, particularly in the presence of free chloride.A kinetic study dealing with the oxidation of [PtCl 4 ] 2− by hydrogen peroxide in perchloric acid has been reported. 8n the present study, oxidation of [PtCl 4 ] 2− by hydrogen peroxide is revisited, with the added complexity of free chloride ions in solution, presenting a more complex mechanism.Formation of [PtCl 5 (H 2 O)] − and possibly trans-[PtCl 4 (H 2 O) 2 ] coincides with [PtCl 4 ] 2− assisted ligand exchange.Apart from the complications associated with such reactions, hydrochloric acid catalyses the decomposition of hydrogen peroxide to form hypochlorous acid according to eqn (2).
Oxidation by hydrogen peroxide and hypochlorous acid may therefore coincide to yield two parallel reactions k H 2 O 2 and k HOCl as depicted in the scheme in (3), where the intermediates are Oxidation of [PtCl 4 ] 2− by hydrogen peroxide is slow, whereas oxidation by hypochlorous acid is rapid in comparison, but yields the same product. 8,9Furthermore, depending on the reaction conditions, hypochlorous acid may form chlorine which can also act as an oxidant, further complicating this process.These aspects are explored and discussed in the present study to elucidate the overall oxidation mechanism of [PtCl 4 ] 2− by hydrogen peroxide in the presence of an excess of hydrogen cations and chloride anions.

Spectrophotometry
Photo-induced reactions necessitate kinetic measurements in UV-Vis absorption spectra at a specific wavelength.UV-Vis spectra in the range 200-600 nm were recorded after the oxidation was complete.Measurements were performed on a Shimadzu UV-2010PC spectrophotometer.The instrument was equipped with a thermoelectrically controlled cell holder using 1 cm tandem quartz cuvettes.Activation volume measurements for slow reactions were performed on a Shimadzu UV-2010PC spectrophotometer equipped with a high pressure cell fitted with a 1.5 cm pill-box quartz-cuvette. 14Activation volume measurements for relatively fast reactions were performed on a laboratory-made high-pressure stopped-flow instrument. 15The temperature was controlled and maintained in these instruments at 35.0 ± 0.1 °C using a circulating water bath (Julabo MP-5).

Kinetics
Observed rate constants for pseudo-zero order reactions were obtained directly from the slope of concentration versus time plots.Observed rate constants for reactions showing pseudofirst order character were calculated directly from absorbance versus time plots using a least-squares program.The ionic strength was kept constant at 1 M in all experiments by using the correct ratios of hydrogen chloride, sodium chloride, sodium perchlorate and perchloric acid.Hydrogen peroxide was always present in large excess (>15 times) with regard to the substrate ensuring pseudo-order reaction conditions.Small differences in absorbance versus time plots collected at a specific wavelength were noticed when compared to spectra collected in the wavelength range 200-500 nm on a rapid-scan UV-Vis spectrophotometer.It has long been known that Pt(II/IV) complexes are sensitive to light resulting in photoinduced reactions. 10,11The rates of [PtCl 4 ] 2− and/or [PtCl 6 ] 2− aquation as well as [PtCl 5 (H 2 O)] − chloride anion reactions are enhanced when spectra are recorded in the 200-500 nm wavelength range, which accounts for the small differences observed in the kinetic traces.To suppress such effects, changes in absorbance were monitored at 262 or 353 nm, close to the absorbance maxima of [PtCl 6 ] 2− .

Spectra
Oxidation of [Pt(CN) 4 ] 2− by chlorine yields trans-[Pt(CN) 4 Cl-(H 2 O)] − as the primary product even when the free chloride ion concentration in solution is as high as 1 M. 9 A preference for solvent water to coordinate the negatively charged Pt(II) precursor complex, trans to the oxidant, was considered here.The spectrum of trans-[PtCl 4 (H 2 O) 2 ] is known and has been reported before. 2 No evidence could be found for the presence of significant concentrations of trans-[PtCl (5-100 mM) generate predominantly linear absorbance vs. time plots at 262 nm (Fig. 1; data summarized in Table 1), indicating that the reaction is pseudo-zero order with respect to Pt(II) and the observed rate is defined by rate law eqn (4).
The observed rate constants (k 0 obs ) were obtained directly from the slope of concentration versus time plots.Plots of k 0     1, from which it follows that k 0 = (8 ± 2) × 10 −7 s −1 over the entire concentration range investigated.
The activation parameters ΔH ‡ and ΔS ‡ were estimated from the variation of k 0 obs in the temperature range 15-35 °C (Table 2), and plots of ln (k 0 /T ) versus 1/T according to the Eyring equation gave linear correlations (ESI Fig. S2 †).Values for ΔH ‡ and ΔS ‡ were subsequently calculated from the slope and intercept of such plots according to eqn ( 6) and (7), where k B is Boltzmann's constant and h is Planck's constant.The calculated parameters are listed in Table 2.
The activation volume of the oxidation reaction (ΔV ‡ ) for the zero-order process was estimated by varying the pressure in the range 10-132 MPa.A plot of ln k 0 versus pressure is linear (Fig. S3, ESI †) and ΔV ‡ was obtained from the slope of the plot according to eqn (8).The calculated activation parameters are included in Table 3.
Oxidation  S5 †), suggesting that the reaction is first order with respect to hydrogen peroxide.The second order rate constants k H 2 O 2 were calculated from eqn (10) to give the average value (1.5 ± 0.1) × 10 Oxidation of 1 mM [PtCl 4 ] 2− with 300 mM H 2 O 2 generated traces illustrating pseudo-first order character.Activation parameters were estimated under these reaction conditions from the variation of k H 2 O 2 in the temperature range 15-35 °C and eqn ( 6) and (7) (Table 2).Values of ΔV ‡ were estimated by varying the pressure between 5 and 152 MPa for the oxidation of 1 mM [PtCl 4 ] 2− with 15 mM H 2 O 2 .A plot of ln k H 2 O 2 versus pressure is linear (Fig. S6, ESI †) and allowed the estimation of ΔV ‡ from eqn (8) listed in Table 3.

Effect of acid and chloride concentrations on the oxidation rate and reaction order
The effect of variation of acid and/or chloride concentration on the reaction rate and order was evaluated in the ranges 0.6-1 M and 0.02-1 M, respectively.Concentrations of Pt(II) The general rate law eqn (13) was established to define the overall rate of decomposition via the parallel reactions eqn (2) and (12).
Reaction ( 12) is slow relative to the acid-dependent decomposition of hydrogen peroxide, reaction (2), especially in an acidic matrix which will enhance the acid-dependent reaction.The second term in eqn (13) becomes negligible in 1 M H + , as used in this study, so that the rate law for decomposition of hydrogen peroxide reduces to eqn (14).
Zero-order reaction mechanism If a [PtCl 4 ] 2− solution is mixed with excess hydrogen peroxide and hydrochloric acid, hydrogen peroxide is consumed via reactions denoted by k H 2 O 2 and k HOCl in the scheme outlined in eqn (3).Since oxidation of [PtCl 4 ] 2− with hypochlorous acid is rapid, the k 1 reaction is expected to be the major oxidation reaction under the condition of eqn (15). 9,19 The condition is fulfilled where [PtCl 4 ] 2− ≪ 0.5 mM for [Cl − ] = [H + ] = 1 M. ‡ Under such conditions hydrogen peroxide will disappear via the consecutive reactions k 1 → k HOCl depicted in the scheme in (3), where k 1 is the rate-determining reaction for the conditions of eqn (16).

First-order reaction mechanism
2− ] ≥ 0.2 mM in  this work as listed in Table 4 are at the upper limit of this range.A deviation between expected and calculated values may be reasonable, since estimation of k obs values from rate law eqn (9) does not account for any oxidation by hypochlorous acid, the parallel reaction under these reaction conditions.

Rate dependence on chloride and acid concentrations
The oxidation rate and order of [PtCl 4 ] 2− is directly proportional to both chloride and acid concentrations as illustrated by the k obs values shown in Table 5.In order to keep the rate law consistent for the range of reaction conditions investigated, values of k 0 obs were obtained directly from concentration versus time plots.This approach does not account for the parallel first-order oxidation reaction, a possible consequence of the small intercept in these graphs.The overall rate law accounting for both reaction pathways is expressed by eqn (17), and is the sum of eqn ( 4) and (9).
Chlorine as an oxidant Several attempts to determine a rate constant for the oxidation of [PtCl 4 ] 2− by chlorine have failed. 9,20,21Nevertheless, the rapid oxidation of Pt(II) complexes by chlorine is well established.Oxidation of [Pt(CN) 4 ] 2− by chlorine for instance exceeds the oxidation rate of [Pt(CN) 4 ] 2− by hypochlorous acid by five orders of magnitude. 9Hypochlorous acid and chlorine are in rapid equilibrium in aqueous solution according to eqn (18).
Equilibrium ( 18) has been responsible for some of the confusion in the literature, where apparent oxidation by chlorine was in fact oxidation by hypochlorous acid. 9,20,21If the formation of chlorine as a result of eqn ( 18) and subsequent oxidation of [PtCl 4 ] 2− exceeds the rate of formation and oxidation by hypochlorous acid as defined by eqn (19), the overall oxidation reaction will proceed via consecutive reactions depicted by k 1 → k 2 → k Cl2 in the scheme given in eqn (3) and chlorine must be considered as an oxidant in the mechanism discussed here.
Chlorine is the dominant species present in hypochlorous acid/chlorine solutions at equilibrium (> 99.9%, Cl 2 = 1 mM).However, since [PtCl 4 ] 2− is present in large excess relative to the catalytic concentrations of hypochlorous acid formed by eqn (2), chlorine hydrolysis via k −2 according to eqn (19) will be negligibly slow by comparison.The value of k 2 is given as 2.66 × 10 4 M −2 s −1 in the literature. 9Substituting this value into eqn (19) gives the condition that [Pt(II)] ≪ 60 mM when [H + ] and [Cl − ] are 1 M.This is the case in our solutions where [Pt(II)] = 1-20 mM, and implies that Cl 2 must be considered as an oxidant here if we assume that k Cl2 ≥ k H 2 O 2 , which is not unreasonable.Oxidation of [PtCl 4 ] 2− by chlorine will favour formation of [PtCl 6 ] 2− as opposed to the formation of [PtCl 5 (H 2 O)] − when hypochlorous acid is the major oxidant.Since the [PtCl 5 (H 2 O)] − /[PtCl 6 ] 2− ratio increases under conditions that favour the pseudo-zero order mechanism, it supports the idea that hypochlorous acid is the major oxidant and not chlorine.

Conclusions and final comments
The mechanism of oxidation of aqueous [PtCl 4 ] 2− by H 2 O 2 in the presence of excess hydrochloric acid is remarkably complex.Our results obtained show that oxidation takes place at least via two parallel reactions with hypochlorous acid and hydrogen peroxide.The ΔH ‡ and ΔS ‡ values support bond formation prior to electron transfer for oxidation by both H 2 O 2 and HOCl (Table 2).These values are comparable to values reported for the oxidation of [PtCl 4 ] 2− by H 2 O 2 in 1 M HClO 4 (viz.ΔH ‡ = 76 ± 3 kJ mol −1 and ΔS ‡ = −35 ± 9 J K −1 mol −1 ) pointing toward a similar mechanism. 8Activation volumes estimated here for the first-order and zero-order mechanisms are almost identical (Table 3).The negative ΔV ‡ values are characteristic of oxidative addition reactions, i.e.H 2 O 2 -Pt and/or HOCl-Pt bond formation prior to electron transfer, indicating that oxidation takes place via a similar mechanism for both oxidants. 22hese observations are in line with an inner-sphere one-step two-electron transfer mechanism typical of Pt(II) square planar complexes.Since square-planar Pt(II) complexes have a vacant coordination site in the axial plane, formation of an innersphere complex prior to electron transfer seems reasonable.
1.53 × 10 −2 Mean value = (1.5 ± 0.1) × 10 −2 M −1 s −1 Dalton Transactions Paper This journal is © The Royal Society of Chemistry 2014 Dalton Trans., 2014, 43, 6308-6314 | 6311 Open Access Article.Published on 20 December 2013.Downloaded on 4/25/2024 8:28:37 PM.This article is licensed under a Creative Commons Attribution 3.0 Unported Licence.and hydrogen peroxide were kept constant at 0.04 mM and 80 mM, respectively.For the high end concentration range, zero-order kinetics predominated (ESI Fig. S7, ESI, † for the variation of [Cl − ]), while a gradual change to first-order kinetics was observed for sufficiently low concentrations of [H + ] and/or [Cl − ].Rate constants (k 0 obs ) were estimated from the slope of concentration versus time plots.Traces that display a degree of first-order character were evaluated for the first 500 s (Fig. S7, † plots 3 and 4).The values for k 0 obs , obtained in this manner, are listed in Table 5. Plots of k 0 obs versus [H + ] and k 0 obs versus [Cl − ] are linear (R 2 > 0.99) with a small intercept in both cases (Fig. S8, ESI, † −3.4 × 10 −7 M s −1 and 20.3 × 10 −7 M s −1 , respectively), suggesting that k 0 obs varies with [H + ] and [Cl − ] according to eqn (11): k 0 obs ¼ a þ b½Q ð 11Þ where Q represents [H + ] or [Cl − ], a = 3.4 × 10 −7 M s −1 /20.3 × 10 −7 M s −1 and b = 8.6 × 10 −7 s −1 /82.3 × 10 −7 s −1 , respectively.These values were obtained directly from the intercept and gradient of the plots in Fig. S8.† Discussion Rate law for hypochlorous acid formation An oxidation mechanism for [PtCl 4 ] 2− displaying zero-order kinetics can be envisaged as a result of the rapid formation of hypochlorous acid according to eqn (2).Catalytic decomposition of hydrogen peroxide in hydrochloric acid has been studied in detail.It was concluded that hydrogen peroxide decomposes via the acid-dependent reaction eqn (2), in addition to an acid-independent pathway eqn (12).

- 19
If the concentration range of [PtCl 4 ] 2− > 0.5 mM, with [Cl − ] = [H + ] = 1 M, oxidation of [PtCl 4 ] 2− by hydrogen peroxide becomes dominant and the rate of hypochlorous acid formation is small by comparison.Under such reaction conditions, oxidation via the consecutive reactions k 1 → k HOCl competes for the oxidation of [PtCl 4 ] 2− with reaction k H 2 O 2 depicted in the scheme given in eqn (3).In the absence of free chloride ions, oxidation of [PtCl 4 ] 2− by hydrogen peroxide is first-order with respect to both [PtCl 4 ] 2− and H 2 O 2 according to the rate law in eqn (

Table 2
Kinetic data for the oxidation of [PtCl 4 ] 2− by H 2 O 2 as a function of temperature with calculated activation parameters

Table 3
Kinetic data for the oxidation of [PtCl 4 ] 2− by H 2 O 2 as a function of pressure with calculated activation volumes at 35 °C

Table 4
Experimental conditions and kinetic data for the oxidation of 0.2-1.0 mM [PtCl 4 ] 2− yielding pseudo-first order kinetics at 35 °C

Table 5
Oxidation of [PtCl 4 ] 2− performed to determine the effect of chloride and acid concentrations on the observed reaction order at 35 °C Calculated from eqn (15) with k H 2 O 2