Clarifying the Equilibrium Speciation of Periodate Ions in Aqueous

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Clarifying the Equilibrium Speciation of Periodate Ions in Aqueous Medium László Valkai,† Gábor Peintler,‡ and Attila K. Horváth*,† †

Faculty of Sciences, Department of Inorganic Chemistry, University of Pécs, Pécs, Hungary Faculty of Science and Informatics, Department of Physical Chemistry and Material Sciences, University of Szeged, Szeged, Hungary

‡

S Supporting Information *

ABSTRACT: Equilibria of periodate ion were reinvestigated in aqueous solution by using potentiometric titration, UV and Raman spectroscopies, and gravimetry simultaneously at 0.5 M ionic strength and at 25.0 ± 0.2 °C. Stepwise acid dissociation constants of orthoperiodic acid were found to be pK1 = 0.98 ± 0.18, pK2 = 7.42 ± 0.03, and pK3 = 10.99 ± 0.02, as well as pK2 = 7.55 ± 0.04 and pK3 = 11.25 ± 0.03 in the presence of sodium nitrate and sodium perchlorate as background salts, respectively. pK1 cannot be determined unambiguously from our experiments in the presence of sodium perchlorate. The molar absorptivity spectrum of H4IO−6 and H3IO2− 6 was determined in the range of 215− 335 nm, as major species of periodate present from slightly acidic to slightly alkaline conditions. The solubility of periodate decreases significantly under alkaline conditions, and it was determined to be (2.8 ± 0.4) mM by gravimetry, under our experimental conditions. None of these studies gave any clear evidence for an ortho−meta equilibrium and the frequently invoked dimerization of periodate. All measurements can quantitatively be described by the presence of orthoperiodic acid and its three successive deprotonation steps.

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INTRODUCTION Periodate is well-known as a strong oxidizing agent with various chemical applications. For example, a step of a cyclocondensation reaction requires periodate as a coreactant to synthesize large cyclic systems,1 to produce cyclic dinucleotide analogues2 and new cellulose-composite materials.3 A wide variety of its application covers C−H-bond functionalization in hydrocarbons, oxidative iodination of arenes, sulfonylation of aromatic compounds, aziridination, azidoiodination, hydroxyhalogenation, dihydroxylation, diazidation of styrene derivatives, epoxidation; oxidation of alcohols, α-oxidation of carbonyl compounds, oxidation of sulfides, selenides, thioureas, selenoureas, indoles, tetrahydro-β-carbolines, and aromatic amines, and even oxidative rearrangements may also be found.4 Surprisingly, as a result of the synergism of periodate and osmium tetroxide, even oxidation of alkanes may be driven at relatively mild conditions.5 Nowadays, there is still no clear explanation for this synergistic effect. Periodate may be found in numerous solid forms: sodium metaperiodate (NaIO4), trisodium dihydrogen orthoperiodate (Na3H2IO6, also named as sodium triparaperiodate), barium periodate (Ba(IO4)2), potassium metaperiodate (KIO4), and orthoperiodic acid (H5IO6). In the 1950s, Crouthamel and his co-workers reported detailed equilibrium studies to determine the existing species in aqueous medium and to calculate the equilibrium constants of the corresponding processes.6,7 Three deprotonation steps of orthoperiodic acid form were found with the values of pK1 = 1.64, pK2 = 8.36, and pK3 ≈ 14.98 at low ionic strength in the presence of perchlorate as a background electrolyte. © XXXX American Chemical Society

A bit later Siebert reported the Raman spectra of periodate solution in 1953.8 In his work, he claimed an experimental evidence for an ortho−meta equilibrium of pK

H5IO6 HooI IO−4 + H+ + 2H 2O

which is able to explain the determined deprotonation constant of pK = 2.3 × 10−2. Four Raman peaks at 791 (A1), 853 (F2), 256 (E), and 325 cm−1 (F2) were reported for the ortho form of periodate, and an “extra” line at 636 cm−1 was assigned to the meta form of periodate. In 1968, further proof of the ortho−meta equilibrium was reported by Kren et al.9 They used NMR spectroscopy to investigate the hydrolytic process. Two parameters of the derivative of the periodate peak were used: the height of the maximum and the half of the full width at half-maximum with mean average deviation of 4% and 2%, respectively. The equilibrium constant for the H4IO−6 ⇌ IO−4 + 2H2O water elimination reaction was determined to be 29 ± 2 at 25 °C. In contrast to this, Kerezsi and co-workers also investigated this equilibrium by NMR spectroscopy and found no evidence of the described process.10 One of the main outcomes of their study may be summarized that in aqueous periodate solution the tetrahedral form periodate is significantly less dominant than previously thought. In 1965, Buist and Lewis reported a possible dimerization process of periodate.11 They claimed that pK2 (same as above) Received: July 28, 2017

A

DOI: 10.1021/acs.inorgchem.7b01911 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

preparation and the measurement the solutions were stored at dark. The temperature in every measurement was kept constant at 25 °C. The ionic strength was adjusted by sodium perchlorate to 0.5 M, except otherwise stated. During titrations, all solutions were stirred at a constant 500 rpm. Nitrogen flow was maintained above the solution to avoid gas dissolution from the air. The base solution was dosed by a conventional 25.00 mL borosilicate glass buret equipped with solid glass plug. Gravimetric Analysis. Sodium metaperiodate (0.10−0.92 g) was dissolved in ∼35 mL of water, and 5 mL of 5 M sodium perchlorate solution was added to it. In the concentration range applied, the solutions were always clear, no precipitation, and not even opalescence was noticed. Then 0.8 mL of concentrated aqueous sodium hydroxide solution (resulting in pH higher than 12) was added, and the volume was set to 50.00 mL by distilled water. The solutions were kept in dark and were homogenized twice a day during the next 11 d. After this aging period, the precipitate was removed by glass filters. Before the filtration ultrasonic bath was used for 2 min to remove the solid particles from the wall of the volumetric flasks. Portions of the mother liquor (each ∼15 mL) were used three times to collect all the solid material. The solid precipitation was then washed by 10 mL of distilled water without using vacuum to remove the alkaline solution from the surface of the samples. Then vacuum was turned on again for 1 min to dry the sample. The solid dry samples were then dried on 66 °C for 12 h followed by a natural cooling period for 7 h at room temperature. Potentiometry. Two sets of pH-metric analysis were performed. In all cases, 33 mM sodium hydroxide solution having an ionic strength of 0.5 M is added to a stirred 10 mM acid solution containing the measured analytes, such as sodium periodate or succinic acid. In the first set of experiments nitric acid was used, and the ionic strength was adjusted by sodium nitrate. In the second set, perchloric acid was used as a strong acid, and the ionic strength was adjusted by sodium perchlorate. After a two-point calibration procedure, a strong acid− strong base and a succinic acid with added strong acid−strong base titrations were performed. These titrations were necessary for the accurate calibration of the instrument with the pHCali program24 to determine the Irving factor,25 the slope of the Nernst equation, the water ionization constant, the concentration of the strong base, and the incidentally dissolved carbon dioxide amount in the system. These parameters are then used to convert the measured pH value of a solution containing periodate into equilibrium concentration of the proton that allows us to fit the acid dissociation constant(s) of the periodate at the ionic strength given. UV Spectroscopy. UV spectra of a 19 mM phosphoric acid solution titrated against 32 mM sodium hydroxide solutioneach of it contained 93.5 μM sodium periodatewere registered in a quartz cuvette having an optical path of 1.000 cm. During the titration process, the pH values of the solutions were also registered simultaneously. The whole process was repeated without using periodate to measure the absorbance of the background electrolyte. Ionic strength of these solutions was mainly controlled by the buffer components used; no extra background salt was applied to avoid its overwhelming contribution to the measured absorbance. Only absorbance lower than 0.95 was used for data evaluation. Raman Spectroscopy. The Raman spectra of solid materials was collected from 64 scans, the spectra of the solutions measured in a quartz cuvette were averaged from 1024 scans in each case, and the spectral resolution was 4 cm−1. Instruments. Analytik Jena SPECORD S600 diode-array spectrophotometer with thermostated cell holder was used to register the UV/vis spectra. Thermo Nicolet 5700 FT-IR spectrometer with Raman accessory, with a 980 nm laser source, was used to record the Raman spectra. Sherwood 410 Flame photometer was applied to measure the sodium content of our samples. The pH values were followed by a Boeco Germany pH meter, attached to a combined pH electrode (BOE 5095626) filled with an appropriate liquid electrolyte. Data Treatment. Exporting the ASCII files was done with the software of each instrument. Raman spectra were treated and evaluated by Spectragryph.26 Matrix rank analysis was performed to determine the number of independent absorbing species from the UV-spectra by

is dependent not only on the temperature but also on the total concentration of periodate, as well. They used periodate in the 2−50 mM concentration range and reported changing pK2 values at 0.1 M ionic strength with respect to the overall periodate concentration at 0, 25, and 45 °C. Four years later they try to find further support for the dimerization by UV, IR, and Raman spectroscopies.12 Originally, Buist and Lewis tried to find a plausible explanation to interpret the varying formal kinetic order of periodate in the pinacol−periodate reaction. For this reason they introduced the dimerization of periodate as a feasible possibility. Several years later, however, Buist et al.13 offered an even more reasonable explanation for the same phenomenon, namely, the accumulation of an intermediate. Consequently, this subsequent paper questioned their own previous suggestion about the existence of dimerization. Surprisingly, this conclusion has fallen into oblivion, and previous erroneous interpretation has been generally accepted. The most complex model about the system was published in 1980 by Kustin and Simoyi.14 They reported the deprotonation, dimerization, and ortho−meta equilibrium constants. The temperature-jump study applied allowed them to determine the rate coefficients of the forward and backward processes as well. Furthermore, poor solubility of periodate was also reported, based on formation of white crystalline insoluble material upon addition of strong base to the periodate solution. This low solubility reported by Kustin and Simoyi14 seems to contradict the relatively high concentration of periodate used by Buist at alkaline conditions for determination of the dimerization process.11,12 Furthermore, obtaining reliable Raman spectra in aqueous solution requires relatively concentrated periodate solutions that cannot be incorporated with the low solubility found by Kustin and Simoyi.14 The previous literature de facto accepted the existence of both the ortho−meta equilibrium as well as the dimerization, and these processes have frequently been invoked.15−17 There are many interests for different reactions of periodate, and our research group investigated some of them quite recently. Reactions with chlorite,18 bromide,19 thiosulfate,20 tetrathionate,21 iodide,22 pentathionate23 were soundly described, though during this research some contradictions were noticed between our experiments and the previous literature. Therefore, we decided to perform a complete reinvestigation of the periodate equilibrium in aqueous medium. The present work led to a significantly simpler equilibrium model that agrees best with the earliest one proposed by Crouthamel et al.6,7 including successive deprotonation of periodate ion and does not support any further and unnecessary complication caused by any dimerization process or the involvement of ortho−meta equilibrium.

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EXPERIMENTAL SECTION

Chemicals. Sodium perchlorate monohydrate (Sigma-Aldrich), potassium nitrate (Reanal), potassium chloride (Sigma-Aldrich), sodium chloride (Reanal), sodium carbonate (Reanal), sodium hydrogen carbonate (Reanal), sodium metaperiodate (Reanal), nitric acid (VWR), absolute ethanol (VWR), methyl red (Reanal), and succinic acid (Reanal) were of the highest purity commercially available and used without further purification. The twice ionexchanged water was double-distilled atmospherically to avoid any impurities of the exchange resin. Stock solutions were prepared by dissolving the calculated amount of the chemicals. Strong acid stock solutions were stored at dark in a closed vessel, and the concentration was monitored by classical acid−base titrations. All the other solutions were prepared freshly and used the day of preparation. Between the B


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Inorganic Chemistry MRA program package.27 Titration data were evaluated by pHCali24 and PSEQUAD28 programs. Deconvolution of the UV spectra was performed by ChemMech29 program package.

Table 1. Conceivable Sodium-Containing Speciesand Their Sodium Content Expressed in Mass Percentin Aqueous Periodate Solution

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RESULTS AND DISCUSSION Gravimetric Analysis. Sodium metaperiodate is well soluble in water; even a solution having a concentration of 0.4 M can be prepared. White precipitation occurs, however, upon addition of sodium hydroxide. This fact was noticed by Kustin and Simoyi,14 and they suggested that the solid material produced is a dimer of the periodate (Na4H2I2O10). To verify this statement, gravimetric experiments were performed. The mass of the dry precipitate over the weighed sodium metaperiodate is given at the top of Figure 1. The lower part

formula

Mr

wNa

NaH4IO6 Na2H3IO6 Na3H2IO6 Na4HIO6 Na5IO6 NaIO4 Na6I2O10 Na5HI2O10 Na4H2I2O10 Na3H3I2O10 Na2H4I2O10 NaH5I2O10

249.92 271.90 293.88 315.86 337.84 213.89 551.73 529.75 507.77 485.78 463.80 441.82

9% 17% 23% 29% 34% 11% 25% 22% 18% 14% 10% 5%

periodate, although it does not necessarily mean the exclusion of a periodate dimer in the liquid phase. As seen the seventh and eighth species (Na6I2O10 and Na5HI2O10) in Table 1 might be reasonable candidates for being the actual composition of the precipitate. Therefore, potentiometric titration was also performed to seek an evidence for the existence of a dimer in solution. Potentiometry. Two potetiometric titrations were performed at 0.625 and 1.25 mM total periodate concentration using nitrate as a background electrolyte. The results are represented by the black and red curves in Figure 2,

Figure 1. (top) Mass of the dry precipitate (mprec) as the function of mass of the weighted sodium metaperiodate (TNaIO4) in the gravimetric experiments. Every point represents a different solution. (bottom) Raman spectra of each samples. Figure 2. Experimental and calculated pH over the volume of added strong base. Dots represent the measured data; the solid lines are the calculated values after the fitting procedure. The blue and the green curves were measured in perchlorate-containing medium; the others were obtained in sodium nitrate as a background salt.

of the figure shows the Raman spectra of each solid sample providing an evidence that the composition of the solid material was exactly the same in every case. A linear function fits perfectly to the measured points of the mass−mass graph. From the parameters of the fitted straight line, the concentration of the saturated alkaline solution of periodate was calculated to be (60 ± 9) mg/100 mL that corresponds to (2.8 ± 0.4) mM. The composition of the product was determined from its sodium content measured by atomic absorption spectroscopy. The result shows (23 ± 1) wt % sodium in the solid precipitate, from which the Na3H2IO6 formula can be deduced most likely as seen from Table 1. The sodium content and the perfect linear dependency of the mass of the precipitate over the mass of periodate together straightforwardly prove that the solid precipitate cannot be the previously suggested14 dimeric form (Na4H2I2O10) of

respectively. The curvature of the measured pH−volume data suggests more deprotonation steps; therefore, the calculations were performed by the following model (where HA represents the added strong acid): K1

H5IO6 ⇌ H4IO−6 + H+ K2

H4IO−6 ⇌ H3IO62 − + H+ K3

H3IO62 − ⇌ H 2IO36 − + H+ C


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Inorganic Chemistry

Spectrophotometric Studies. Earlier reports6,7,11,12 came to the conclusion that monomer form of periodate has a characteristic UV-absorption peak at 210 nm. Consequently, if a dimer is present at a suitable pH range, one may expect a significant red shift to lower energy levels in its absorption maximum. Even though the concentration of periodate must be kept at a low level to measure meaningful absorbances, the change in UV spectra should follow this feasible expectation. The spectra obtained are shown in the Figure 3. During the titration process in the acidic region, a peak having a maximum

H 2O ⇌ H+ + OH− H 2CO3 ⇌ H+ + HCO−3

HCO−3 ⇌ H+ + CO32 − HA → H+ + A−

Since it is impossible to prepare completely carbonate-free sodium hydroxide solution, the slight pH effect of these equilibria was also taken into consideration. Note, however, that the carbonate contamination never exceeded 0.5% of the total base concentration. As a result of the simultaneous evaluation of pH−volume titration curves, the following acid dissociation constants can be calculated for the stepwise equilibria of orthoperiodic acid: pK1 = 0.98 ± 0.18, pK2 = 7.42 ± 0.03, and pK3 = 10.99 ± 0.02. The average deviation between the measured and calculated pH was found to be 0.006 pH unit indicating a very sound description of the experimental data. The value of pK2 nicely coincides with the one determined by Crouthamel and his co-workers,6,7 but there is a 0.7 pH unit difference for pK1. Furthermore, pK3 shows a huge difference compared to the value determined by Crouthamel et al.7 Therefore, we also repeated our experiments in a different medium, where sodium perchlorate was used as a background electrolyte. In the next set of experiments, potassium chloride electrolyte of the electrode was changed to sodium chloride solution to avoid the formation of potassium perchlorate precipitate. The concentration of periodate was increased to 1.0 and 2.0 mM. The graphical representation of the fit is also shown in Figure 2, by blue and green colors, respectively. The calculations resulted in the pK2 = 7.55 ± 0.04 and pK3 = 11.25 ± 0.03 values with an average deviation of 0.007 pH unit. From these experimental data, pK1 cannot be determined within acceptable standard deviation. Comparing these values to the ones determined in nitrate medium, slightly higher pK values were found, but these slight differences can easily be explained by the medium effect of different background electrolytes. Reproducibility of the profile of titration curves was also checked by titrating the periodate solution first with a strong base followed by a back-titration of this alkaline solution by a strong acid. Figure S2 of the Supporting Information clearly proves a perfect reproducibility indicating that reversible and relatively rapid processes must be taken into consideration for describing the result. To study the possible pH effect of the dimerization process, further simulations were performed by PSEQUAD.28 Our original model was therefore extended by the following equilibria:

Figure 3. Measured UV spectra (circumstances are given in the Experimental Section). The first spectrum (light green) is recorded in the initial phosphoric acid solution; the next ones (red) are measured in acidic conditions and in alkaline solutions (blue).

at λmax ≈ 222 nm is observed whose intensity increases by an increase of pH; meanwhile, an isosbestic point appears at ∼205 nm. As pH increases in alkaline medium the spectra become more complex, as two isosbestic points appear at 213 nm and at 245 nm, simultaneously. There are straightforward connections between the number of isosbestic points and the electron transition peaks; thus, we assume three different light-absorbing species: the first peak must belong to periodic acid (H5IO6), a major species at low pH values. As pH increases we can see the appearance of an additional peak corresponding to H4IO−6 species. Finally, under alkaline conditions the curvature of UV spectra along with the two isosbestic points might perhaps suggest the appearance of two independent electron transition peaks. To determine the number of linearly independent absorbing species present in the system, matrix rank analysis was performed by the MRA program package.27 This program is able to calculate residual absorbance as a function of the wavelength supposing different number of linearly independent absorbing species. During the calculation procedure 24 recorded spectra (210−400 nm) were used in the pH range between 2.0 and 8.9. The result is visualized in Figure 4, from which it is unambiguously concluded that the system can be described by two absorbing species. To extract reasonable information about the absorption bands of the different forms of periodate from the UV spectra, deconvolution was performed. The wavelengths used were larger than 194.5 nm, higher absorbances than 0.95 were neglected (where linearity of absorbance as a function of concentration is not strictly fulfilled), and all the 31 measured spectra were used for data evaluation. As a first step, wavelengths were transformed into wavenumbers, because the

5− 5− H3IO62 − + H 2IO36 − ⇌ H5I 2O12 ⇌ HI 2O10 + 2H 2O 4− 4− 2H3IO62 − ⇌ H6I 2O12 ⇌ H 2I 2O10 + 2H 2O

The calculations revealed that dimerization can cause significant pH change only if more than ∼70% of periodate forms a dimer (see Figure S1 of the Supporting Information), although the change can be more pH units in case of nearly stoichiometric dimer formation. Thus, this method is not so sensitive to support the partial existence of any dimer form, but the given experiments exclude the possibility of stoichiometric dimerization. D


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Inorganic Chemistry

Figure 4. Graphical representation of the result of the matrix rank analysis. Number of assumed, linearly independent light-absorbing species was changed: zero species (blue), one species (red), two species (purple), and three species (green).

latter physical property is proportional to the energy. Then the following Gaussian curve was fitted to the measured absorbance−wavenumber functions: A n −(ν ̃− νmax ̃ )2 /2σ 2 f (ν ;̃ A n , νmax e ̃ , σ) = σ 2π where ν̃ is the wavenumber (i.e., the independent variable), and ν̃max, σ, and An are the adjustable parameters, that is, the wavenumber at the peak maximum, the deviation related to the half-width of the peak, and a scaling factor, respectively. During the course of calculation procedure, one ν̃max value and one σ value were fitted for the same peaks of each spectrum, while the scaling factor was determined individually for each spectrum and at every electronic transition. Performing the calculation process maximum of 198 fitted parameters was used to describe the measured data simultaneously as shown in Table 2. The number of fitted parameters, of course, depends on the number of peaks fitted.

Figure 5. Results of the deconvolution of the measured spectra at different pHs. (top) Result of a typical deconvolution (pH = 8.5), each curve represents one electron-transition band. (bottom) Fitted parameter An vs the pH. (inset) After shifting the two blue curves to a common baseline followed by rescaling them the overlapping of these curves is visualized.

the scaling factors versus the measured pH. As seen the latter one indicates a remarkable resemblance on a distribution curve. Further results of the deconvolution at different pHs are also illustrated in the Supporting Information (see Figure S3). The fitting procedure provided the following places for the absorption peak maxima: λmax,1 = 209.7 nm (sky blue), λmax,2 = 193.4 nm (orange), λmax,3 = 223.5 nm (green), and λmax,4 = 229.6 nm (dark blue). The green curve at 223.5 nm belongs quite likely to the form of H4IO−6 as the contribution of this species to the measured absorbance basically disappears above pH = 9.0. The orange curve may be assigned to the absorption of the deprotonated buffer component, since its contribution increases toward higher pHs. Interestingly, the two blue curves (having absorption maxima at 210 and 230 nm) appear simultaneously around pH = 7, and they seem to be dependent on each other, as the inset of Figure 5 suggests. Two feasible explanations may be given to encounter this observation. One option might be if the peak having an absorption maximum at 210 nm is assigned to H3IO2− 6 and if the one having an absorption maximum at 230 nm is assigned to the dimer form. It would consequently mean that the dimer must form from H3IO2− 6 species and that its formation cannot be accompanied by a change in proton concentration. Additionally, the opposite trend of the blue curves (light and dark blue) in the Figure 5 above pH = 11 might support the dimer formation. This trend may, however,

Table 2. Fitted Parameters during the Deconvolution shared parameters

ν̃max,1

ν̃max,2

ν̃max,3

ν̃max,4

ν̃max,5

ν̃max,6

shared parameters

σ1

σ2

σ3

σ4

σ5

σ6

spectra 1 spectra 2 ⋮ spectra 31

A1 A7 ⋮ A181

A2 A8 ⋮ A182

A3 A9 ⋮ A183

A4 A10 ⋮ A184

A5 A11 ⋮ A185

A6 A12 ⋮ A186

All the parameters were refined to minimize the average deviation between the measured and calculated absorbances. The calculation processes resulted in the following relative average deviations: 2.57%, 2.79%, 0.40%, 0.14%, and 6.12% assuming 2, 3, 4, 5, and 6 peaks, respectively. Conversion of the relative average deviations into absorbance unit leads to 0.022, 0.024, 0.003, 0.001, and 0.052 values. It can clearly be seen that mathematically the best agreement can be achieved if five peaks are assumed, since supposing six peaks leads to seriously overdetermined mathematical model. Taking also into consideration that the uncertainty of absorbance measurement is at least 0.003 au, four Gaussian peaks must be enough to describe simultaneously the measured spectra within the error of the absorbance measurement. The result of a typical fit is shown in the top of Figure 5, while the bottom part indicates E


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Inorganic Chemistry be explained by the next deprotonation step supported by the results of our potentiometric titrations as well. Therefore, in accordance with our potentiometric data the first option to assign the peaks at 210 and 230 nm to H3IO2− 6 and the dimer form of periodate, respectively, may not be appropriate but cannot be ruled out directly. The second conceivable option is that both peaks belong to the same species H3IO2− 6 . Although between these possibilities, no unambiguous decision can be made, following the rule of Occam’s razor we suggest that the peaks mentioned above belong to H3IO2− 6 . The pK2 value and the molar absorptivities of H4IO−6 and H3IO2− 6 were also calculated for both species at every 5 nm in the range of 210−335 nm within the pH range of 3.0−9.5. Twenty-five measured spectra were used for data evaluation. The reason for fitting pK2 value was that the ionic strength during these measurements was significantly less (∼0.1 M) than in case of the potentiometric titrations. The fitted molar absorptivities are given in Figure 6, and pK2 was found to be

Table 3. Exact Values of the Molar Absorptivity Spectra λ (nm)

ελH4IO6− (M−1 cm−1)

ελH3IO62− (M−1 cm−1)

215 220 225 230 235 240 245 250 255 260 265 270 275 280 285 290 295 300 305 310 315 320 325 330 335

8278 ± 5 9586 ± 5 9585 ± 5 8474 ± 5 6762 ± 5 4938 ± 5 3363 ± 5 2181 ± 5 1372 ± 5 857 ± 5 553 ± 5 386 ± 5 298 ± 5 251 ± 5 216 ± 5 184 ± 5 148 ± 5 115 ± 5 82 ± 5 53 ± 5 30 ± 5 0±5 0±5 0±5 0±5

7150 ± 9 6019 ± 10 4983 ± 10 4278 ± 10 3899 ± 9 3662 ± 9 3385 ± 9 2983 ± 9 2469 ± 9 1917 ± 9 1410 ± 9 1003 ± 9 710 ± 9 518 ± 9 393 ± 9 314 ± 9 254 ± 9 209 ± 9 168 ± 9 134 ± 9 103 ± 9 84 ± 9 54 ± 9 35 ± 9 16 ± 8

Figure 6. Fitted molar absorptivity spectra of H4IO−6 and H3IO2− 6 .

8.205 ± 0.003. Significant difference between this value and the ones determined by potentiometric titration may reasonably be explained by the lower ionic strength and the quality of the background electrolyte (phosphate) applied. The values of the molar absorptivities are also given in Table 3. To prove that the relatively huge change in the value of pK2 is mainly the consequence of the medium effect and the lower ionic strength, the photometric titration process was repeated in perchlorate medium at 0.5 M ionic strength. Since no additional buffer components were used, the pH is regulated by the self-buffering effect of periodate. Thus, the pH was adjusted by fractional dropwise addition of sodium hydroxide to reach a pH in the range from 3.0 to 9.0 during the course of titration, and the pH was directly measured by a pH meter calibrated by pHCali method as described above. The deconvolution process was repeated as mentioned previously in this subsection, and the result is illustrated in Figure 7. In this case Arel n denotes the normalized values of parameter An. As it is seen, the equivalent concentration of H4IO−6 and H3IO2− 6 is in complete agreement with the results of the titration and indicates that pK2 must be ∼7.5 by using perchlorate as a medium. This result also confirms that the medium applied may seriously influence the pK2 of orthoperiodate, but it should be emphasized that this effect may not necessarily refer to the existence of a dimer as concluded by Buist and Lewis.11 At the same time it also suggests the presence of phosphoric acid has a relatively huge

Figure 7. Result of the deconvolution in the second measurement. Arel n denotes normalized An values. The matching colors represents the same peaks as above: green corresponds to the H4IO−6 form; the two blue ones belong to H3IO2− 6 .

influence on pK2 of orthoperiodic acid. As generally known pK is dependent on the applied media, and the second protonation step of phosphoric acid is very sensitive to the ionic strength. These two facts may cause the presented change in pK2. Furthermore, this fact is also supported by the IUPAC database,30 which presents several analogous examples in other systems. A word is also in order here to discuss the possible role of the ortho−meta equilibrium. If it exists it should result in two Gaussian curves, and the width at half-maximum of the peaks or the maximum of the absorbances must differ from each other. But only one peak can be detected in the UV region suggesting that the equilibrium is almost stoichiometrically shifted toward the HxIO(5−x)− forms. Siebert proposed that the ortho−meta 6 F


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Inorganic Chemistry

at a wide concentration range from 1 × 10−5 to 1 × 10−2 M, where significant red shift of the absorption band was observed.11 Appearance of the new band was attributed to the formation of a dimer having a dimerization constant of 2400 at pH = 11.4, but it is not clear how the pH was kept constant in their experiments.11 Certainly, innocent buffer, controlling the pH around this region and simultaneously not contributing to the measured absorbance sign, is quite difficultif it can be at allto apply. To our point of view the most probable explanation of the red shift of absorption peak in their experiments is simply due to the change in the equilibrium speciation between differently protonated forms of orthoperiodate caused by unintentional shift of pH upon increasing the initial periodate concentration. More definite statement is difficult to say without knowing the intimate details of experiments done by Buist et al.11 Later, they provided a further proof of dimerization by measuring UV, Raman, and IR spectra.12 The result of their subsequent UV measurements, however, seems to be questionable simply because at pH higher than 9 solubility of periodate is significantly less than 0.1 M, the highest concentration of periodate they used. This fact is also supported by Kustin and Simoyi,14 who reported white precipitation in the concentration range of 5−12 mM of periodate in the pH range of 9.1−9.6 at an ionic strength of 0.25 M. This observation is also in accordance with our present result obtaining the solubility of periodate 2.8 mM in alkaline conditions. One of the main problems of Kustin and Simoyi’s work is, however, that they only supposed but did not prove the formula of the white precipitate (Na4H2I2O10). Our gravimetric result straightforwardly disproves this possibility and supports that rather the low solubility of Na3H2IO6 is responsible for the precipitation. Despite the fact that ortho−meta equilibrium has long been established no persuasive experimental evidence was found. Siebert reported the Raman spectra of the solution8 for the first time from which he postulated the different forms of periodate. Later, Crouthamel et al. reported that titration of periodate in methanol−water mixture results just in a simple sigmoidal curve, and this observation was interpreted as the disappearance of the ortho form.7 Although this assumption seems to be plausible 85:15 v/v% methanol−water mixture is required to transform the titration curve into a simple sigmoidal one suggesting just one protonation step. However, application of this solvent mixture usually shifts significantly the successive acid dissociation constants of a multiprotic acid studied and the water ionization constant by even such a way that only one dissociation constant can be determined. The literature provides several known examples.31−36 Furthermore, if the ortho−meta equilibrium really exists then some spectral evidence should be found. Kren and his co-workers9 used 127 I-NMR spectroscopy to investigate the ortho−meta equilibrium, and they found just one peak and used its change as a function of temperature to calculate the equilibrium constant of the water elimination process of orthoperiodate. Despite the fact that water exchange processes are usually faster than the recorded relaxation times in aqueous solution, precise quantitative information meets difficulties to be extractedif it can be at allfrom the measured spectra. Furthermore, the fact that just one peak could be registered allows us to be skeptical about the existence of two different forms of periodate. This skepticism is further supported by our spectroscopic measurements, since at acidic conditions a simple distribution curve

equilibrium can be verified by Raman spectroscopy.8 Therefore, for the sake of completeness Raman spectra of the acidic periodate solution were also investigated. Raman Spectroscopic Measurements. Figure 8 illustrates the Raman spectra of acidic periodate solutions. Each

Figure 8. Raman spectra of different periodate solutions. The concentrations are the following: [NaIO4]0 = 0.4 M, [NaClO4]0 = 0.5 M (black); [NaIO4]0 = 0.4 M, [HClO4]0 = 0.5 M (blue); [NaIO4]0 = 0.4 M, [HClO4]0 = 1.0 M (green).

sample contained 0.4 M sodium periodate with variable perchloric acid and sodium perchlorate concentrations. Relatively high solute concentration is used, because even at these experimental circumstances the more intense Raman bands may only be detected. Four peaks can clearly be distinguished. The first one has a maximum at 937 cm−1 that belongs to perchlorate ion. At low pH values a band appears at 634 cm−1 that can reasonably be assigned to the protonated form (H5IO6). As pH increases we may find two additional peaks having maxima at 794 and 858 cm−1 and varying simultaneously. These bands can be assigned to the symmetrical and asymmetrical stretching of H4IO−6 ion, respectively. This assignation is a natural consequence of the facts that (a) an asymmetric stretching occurs at slightly higher Raman shifts than the symmetric ones and (b) a general observation that symmetric signs usually have much stronger Raman intensities, compared to the asymmetric ones. On the basis of these facts no evidence was found to support the existence of metaperiodate ion in highly concentrated aqueous solutions, and the assignation of each peak can be easily performed without assumption of existence of this species. Interpretation of the Results. In a good agreement with Crouthamel’s earlier works, stepwise deprotonation constants of orthoperiodic acid were determined in different media and at constant ionic strength. These constants were calculated as pK1 = 0.98 ± 0.18, pK2 = 7.42 ± 0.03, and pK3 = 10.99 ± 0.02, as well as pK2 = 7.55 ± 0.04 and pK3 = 11.25 ± 0.03 in the presence of sodium nitrate and sodium perchlorate media, respectively. Although a slight medium effect may be observed in these cases, the values of pK2 and pK3 indicate a sound coincidence with the ones reported by Crouthamel et al.6,7 These values were later argued by Buist and his co-workers who claimed experimental evidence for the dimer formation of periodate in alkaline medium. Their arguments were based on an indirect evidence of the changing formal kinetic order of periodate in the pinacol−periodate reaction.13 As a result of this observation they measured the UV spectra of periodate solution G


Inorganic Chemistry

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ACKNOWLEDGMENTS The authors are grateful for the financial support of the Hungarian Research Fund NKFIH-OTKA Grant No. K116591. This work was supported by the GINOP-2.3.2-15-2016-00049 grant. The research was supported by the New National Excellence Program of the Ministry of Human Capacities. The present scientific contribution is dedicated to the 650th anniversary of the foundation of the University of Pécs, Hungary.

alone is able to describe soundly the UV as well as the Raman spectra. Finally, also note that the kinetically active form of periodate in aqueous solution was suggested to be the orthoperiodate ion in studying the complex formation of iron(III) hydroxo dimer with periodate ion.10 The most important conclusion of the study by Kerezsi et al. was that, although the rapid interconversion between tetrahedral and octahedral forms of periodate may exist in aqueous solution, previous experimental evidence supporting the existence of this process can easily be interpreted in a different but simpler way as well. According to our results, we completely agree with their statement that new interpretation of published experimental data may easily lead to such a conclusion, where the role of metaperiodate (and its rapid interconversion with orthoperiodate) in aqueous solution is less significant than generally thought.

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CONCLUSION Although studying the equilibria of aqueous periodate solution has long been the subject of intensive research it looked to be completely resolved by the 1980s. It seems to be generally accepted that in aqueous solution periodate is present in a variety of (de)protonated forms of ortho- and metaperiodate ion including the water elimination equilibrium between H4IO−6 and IO−4 and various dimer forms of orthoperiodate ion. Surprisingly, however, neither the ortho−meta nor the dimerization equilibria in aqueous sodium periodate solution can be proven by the use of the same experimental techniques applied previously by other researchers. Our comprehensive study indicates that all the measurements can be interpreted by the successive deprotonation of orthoperiodic acid. Even for the appearance of white precipitation no dimer form of orthoperiodate is needed to be invoked. Our measurements provided an upper limit for the precipitate-free periodate solution to be 2.8 mM, and the most probable composition of white precipitate present under strongly alkaline conditions was found to be Na3H2IO6. The stepwise protonation constants of orthoperiodic acid were also determined by using different background electrolytes. In addition to that the molar absorptivity spectra of the H4IO−6 and the H3IO2− 6 ions were also determined for the first time. Clarifying the speciation of aqueous periodate solution may also contribute to a better understanding of the role of this species in those chemical processes, where complete explanation of the experimental findings is still waiting.5 ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.7b01911. Potentiometry results, UV spectroscopy, numerical data from PSEQUAD simulation (PDF)

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REFERENCES

(1) Hodacova, J.; Budesinsky, M. New Synthetic Path to 2,2,Bipyridine-5,5,-dicarbaldehyde and Its Use in the [3.3] Cyclocondensation with trans-1,2-Diaminocyclohexane. Org. Lett. 2007, 9, 5641−5643. (2) Kinzie, C. R.; Steele, A. D.; Pasciolla, S. M.; Wuest, W. M. Synthesis of Cyclic Dimeric Methyl Morpholinoside - A Common Synthetic Precursor to Cyclic Dinucleotide Analogs. Tetrahedron Lett. 2014, 55, 4966−4968. (3) Codou, A.; Guigo, N.; Heux, L.; Sbirrazzuoli, N. Partial Periodate Oxidation and Thermal Cross-Linking for the Processing of Thermoset All-Cellulose Composites. Compos. Sci. Technol. 2015, 117, 54−61. (4) Sudalai, A.; Khenkin, A.; Neumann, R. Sodium Periodate Mediated Oxidative Transformations in Organic Synthesis. Org. Biomol. Chem. 2015, 13, 4374−4394. (5) Bales, B. C.; Brown, P.; Dehestani, A.; Mayer, J. M. Alkane Oxidation by Osmium Tetroxide. J. Am. Chem. Soc. 2005, 127, 2832− 2833. (6) Crouthamel, C. E.; Meek, H. V.; Martin, D. S.; Banks, C. V. Spectrophotometric Studies of Dilute Aqueous Periodate Solutions. J. Am. Chem. Soc. 1949, 71, 3031−3035. (7) Crouthamel, C. E.; Hayes, A. M.; Martin, D. S. Ionization and Hydration Equilibria of Periodic Acid. J. Am. Chem. Soc. 1951, 73, 82− 87. (8) Siebert, H. Raman-Spektren und Konstitution der Uberjodsaure und des Metaperjodat-Ions. Z. Anorg. Allg. Chem. 1953, 273, 21−27. (9) Kren, R. M.; Dodgen, H. W.; Nyman, C. J. Solubility Product of Cesium Metaperiodate and a Nuclear Magnetic Resonance Study of the Orthoperiodate-Metaperiodate Dehydration Equilibrium. Inorg. Chem. 1968, 7, 446−451. (10) Kerezsi, I.; Lente, G.; Fábián, I. Complex Formation Reaction of the Iron(iii) Hydroxo Dimer with Periodate Ion. Dalton Trans. 2004, 342−346. (11) Buist, G. J.; Lewis, J. D. Dimerization of Periodate in Aqueous Solution, and the Second Dissociation Constant of Periodic Acid. Chem. Commun. 1965, 66−67. (12) Buist, G. J.; Hipperson, W. C. P.; Lewis, J. D. Equilibria in Alkaline Solutions of Periodates. J. Chem. Soc. A 1969, 307−312. (13) Buist, G. J.; Bunton, C. A.; Hipperson, W. C. P. The Mechanism of Oxidation of α-Glycols by Periodic Acid. Part X. The Oxidation of Pinacol, and a General Discussion of the Stability of Periodate Esters and Their Role in the Mechanism of Oxidation. J. Chem. Soc. B 1971, 0, 2128−2142. (14) Kustin, K.; Simoyi, R. H. Kinetics of Orthoperiodate Dimerisation Studied by Temperature-Jump Spectrophotometry. J. Chem. Soc., Dalton Trans. 1980, 2250−2254. (15) Ruff, F.; Fábián, A.; Farkas, O.; Kucsman, A. Mechanism for the Oxidation of Sulfides and Sulfoxides with Periodates: Reactivity of the Oxidizing Species. Eur. J. Org. Chem. 2009, 2009, 2102−2111. (16) Tang, X.; Weavers, L. K. Using Photoactivated Periodate to Decompose TOC from Hydrolysates of Chemical Warfare Agents. J. Photochem. Photobiol., A 2008, 194, 212−219. (17) Tang, X.; Weavers, L. K. Decomposition of Hydrolysates of Chemical Warfare Agents Using Photoactivated Periodate. J. Photochem. Photobiol., A 2007, 187, 311−318.

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AUTHOR INFORMATION

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*E-mail: [email protected]. ORCID

Attila K. Horváth: 0000-0002-1916-2451 Notes

The authors declare no competing financial interest. H


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Inorganic Chemistry (18) Baranyi, N.; Csekő , G.; Valkai, L.; Xu, L.; Horváth, A. K. Kinetics and Mechanism of the Chlorite-Periodate System: Formation of a Short-Lived Key Intermediate OClOIO3 and Its Subsequent Reactions. Inorg. Chem. 2016, 55, 2436−2440. (19) Szél, V.; Csekő , G.; Horváth, A. K. Kinetics and Mechanism of the Oxidation of Bromide by Periodate in Aqueous Acidic Solution. J. Phys. Chem. A 2014, 118, 10713−10719. (20) Rauscher, E.; Csekő , G.; Horváth, A. K. On the Complexity of Kinetics and the Mechanism of the Thiosulfate-Periodate Reaction. Inorg. Chem. 2011, 50, 5793−5802. (21) Filáry, A.; Horváth, A. K. Photochemically Induced Catalysis of Iodide ion and Iodine in the Tetrathionate-Periodate Reaction. Phys. Chem. Chem. Phys. 2010, 12, 6742−6749. (22) Horváth, A. K. Pitfall of an Initial Rate Study: On the Kinetics and Mechanism of the Reaction of Periodate with Iodide Ions in a Slightly Acidic Medium. J. Phys. Chem. A 2007, 111, 890−896. (23) Xu, L.; Horváth, A. K. Autocatalysis-Driven Clock Reaction II: Kinetics of the Pentathionate-Periodate Reaction. J. Phys. Chem. A 2014, 118, 9811−9819. (24) Peintler, G.; Kormányos, B.; Gyurcsik, B. pHCali, A Program for Accurate Calibration of pH-metric Instruments. 2007−2012; http:// www.staff.u-szeged.hu/peintler/, version 1.32a-20070323. (25) Irving, H.; Miles, M.; Pettit, L. A Study of Some Problems in Determining the Stoichiometric Proton Dissociation Constants of Complexes by Potentiometric Titrations Using a Glass Electrode. Anal. Chim. Acta 1967, 38, 475−488. (26) Menges, F. Spectragryph - optical spectroscopy software, Version 1.0.7, 2017. http://www.effemm2.de/spectragryph/. (27) Peintler, G.; Nagypál, I.; Jancsó, A.; Epstein, I.; Kustin, K. Extracting Experimental Information from Large Matrixes. A New Algorithm for the Application of Matrix Rank Analysis. J. Phys. Chem. A 1997, 101, 8013−8020. (28) Zékány, L.; Nagypál, I.; Peintler, G. PSEQUAD for Chemical Equilibria, 1983−2008. 1987−2012; http://www.staff.u-szeged.hu/ peintler/, Version 1.0−5.1. (29) Peintler, G. ChemMech, A Comprehensive Program Package for Fitting Parameters of Chemical Reaction Mechanisms. 1989−2012; http://www.staff.u-szeged.hu/peintler/, Version 2.1−5.0. (30) Martel, E. IUPAC Stability Constants Database. 1997; http:// www.acadsoft.co.uk/scdbase/scdbase.htm, Version 3.09/04. (31) Rochester, C. H. A Comparison of the Acid Ionisation Constants of p-t-Butylphenol, o-t-Butylphenol, and 2,4-di-t-Butylphenol in Water and Methanol. J. Chem. Soc. 1965, 4603−4604. (32) Ohtaki, H.; Tanaka, N. Ionic Equilibriums in Mixed Solvents. VI. Dissociation Constants of Aliphatic Diamines in Water-Methanol Solutions. J. Phys. Chem. 1971, 75, 90−92. (33) Parsons, G. H.; Rochester, C. H. Acid Ionization Constants of 4Substituted Phenols in Methanol + Water Mixtures. J. Chem. Soc., Faraday Trans. 1 1975, 71, 1058−1068. (34) Almasifar, D.; Forghaniha, A.; Khojasteh, Z.; Ghasemi, J.; Sharghi, H.; Shamsipur, M. Spectrophotometric Determination of Acidity Constants of Some Recently Synthesized Anthraquinones in Methanol + Water. J. Chem. Eng. Data 1997, 42, 1212−1215. (35) Psurek, A.; Scriba, G. K. E. Peptide Separations and Dissociation Constants in Nonaqueous Capillary Electrophoresis: Comparison of Methanol and Aqueous Buffers. Electrophoresis 2003, 24, 765−773. (36) Davila, Y.; Sancho, M.; Almandoz, M.; Blanco, S. Structural and Spectroscopic Study of Al(III)-3-hydroxyflavone Complex: Determination of the Stability Constants in Water-Methanol Mixtures. Spectrochim. Acta, Part A 2012, 95, 1−7.

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Clarifying the Equilibrium Speciation of Periodate Ions in Aqueous

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