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1.
Waldemar Gottardi 《Fresenius' Journal of Analytical Chemistry》1998,362(3):263-269
Redox-potential, titrimetric oxidation capacity and pH were used to analyze aqueous iodine solutions over a concentration
range of five decades. By a single operation the equilibrium concentrations of I–, I2, I3
–, I5
–, I6
2–, HOI and OI– were assessed. Accuracy was most influenced by the redox-potential showing a repeatability of < ± 0.1 mV in solutions containing
appreciable iodide (Lugol’s solution) which increased to ±1.5 mV at a 1:100 000 dilution. Together with the uncertainty of
calibration this equals a calculated total error which ranges for “free iodine” ([I2] + [HOI]), from ± 0.5% to ± 4.2% and for total iodide from ± 0.7 to ± 5.3%. Accuracy was assessed by a comparison of the
evaluated equilibrium concentrations and parameters derived therefrom, such as total iodide and redox-potential, with measured
resp. calculated values. The deviations at ionic strength < 0.1 M were below ± 3% and revealed a satisfactory accuracy of
the method. As a unique feature it was possible to monitor the disproportionation in a iodide-free iodine solution with an
initial concentration of 16% hypoiodic acid (HOI) at pH 7.6. Showing a half-life time of only ≈ 75 s it was confirmed that
stability largely depends on HOI concentration.
Received: 26 February 1998 / Accepted: 15 March 1998 相似文献
2.
Studies of the stoichiometry and kinetics of the reaction between hydroxylamine and iodine, previously studied in media below pH 3, have been extended to pH 5.5. The stoichiometry over the pH range 3.4–5.5 is 2NH2OH + 2I2 = N2O + 4I? + H2O + 4H+. Since the reaction is first-order in [I2] + [I3?], the specific rate law, k0, is k0 = (k1 + k2/[H+]) {[NH3OH+]0/(1 + Kp[H+])} {1/(1 + KI[I?])}, where [NH3OH+]0 is total initial hydroxylamine concentration, and k1, k2, Kp, and KI are (6.5 ± 0.6) × 105 M?1 s?1, (5.0 ± 0.5) s?1, 1 × 106 M?1, and 725 M?1, respectively. A mechanism taking into account unprotonated hydroxylamine (NH2OH) and molecular iodine (I2) as reactive species, with intermediates NH2OI2?, HNO, NH2O, and I2?, is proposed. 相似文献
3.
The reduction of iodine by hydroxylamine within the [H+] range 3×10−1–3×10−4 mol.L−1 was first studied until completion of the reaction. In most cases, the concentration of iodine decreased monotonically. However, within a narrow range of reagent concentrations ([NH3OH+]0/[I2]0 ratio below 15, [H+] around 0.1 mol.L−1, and ionic strength around 0.1 mol.L−1), the [I2] and [I3−] vs. time curves showed 2 and 3 extrema, respectively. This peculiar phenomenon is discussed using a 4 reaction scheme (I2+I−⇔︁I3−, 2 I2+NH3OH++H2O→HNO2+4 I−+5 H+, NH3OH++HNO2→N2O+2 H2O+H+, and 2 HNO2+2 I−+2 H+→2 NO+I2+2 H2O). In a flow reactor, sustained oscillations in redox potential were recorded with an extremely long period (around 24 h). The kinetics of the reaction was then investigated in the starting conditions. The proposed rate equation points out a reinforcement of the inhibition by hydrogen ions when [H+] is above 4×10−2 mol.L−1 at 25°C. A mechanism based on ion-transfer reactions is postulated. It involves both NH2OH and NH3OH+ as the reducing reactive species. The additional rate suppression by H+ at low pH would be connected to the existence of H2OI+ in the reactive medium. © 1998 John Wiley & Sons, Inc. Int J Chem Kinet 30: 785–797, 1998 相似文献
4.
The electrochemical behavior of iodine and iodide has been studied in AlCl3+NaCl mixtures with compositions ranging from NaCl saturated melts to AlCl3+NaCl (63+37 mol %) at platinum and tungsten electrodes. Iodide is oxidized in two steps to iodine and I(I); a reduction wave to iodide and an oxidation wave to I(I) are obtained in iodine solutions. The equilibrium constant for the reaction, I?+I(I)=I2, is 6×108 l mol?1 in molten chloroaluminate melts at 175°C. 相似文献
5.
Guy Schmitz 《国际化学动力学杂志》2008,40(10):647-652
An analysis of the former works devoted to the reactions of I(III) in acidic nonbuffered solutions gives new thermodynamic and kinetic information. At low iodide concentrations, the rate law of the reaction IO + I? + 2H+ ? IO2H + IOH is k+B [IO][I?][H+]2 – k?B [IO2H][IOH] with k+B = 4.5 × 103 M?3s?1 and k?B = 240 M?1s?1 at 25°C and zero ionic strength. The rate law of the reaction IO2H + I? + H+ ? 2IOH is k+C [IO2H][I?][H+] – k?C [IOH]2 with k+C = 1.9 × 1010 M?2s?1 and k?C = 25 M?1s?1. These values lead to a Gibbs free energy of IO2H formation of ?95 kJ mol?1. The pKa of iodous acid should be about 6, leading to a Gibbs free energy of IO formation of about ?61 kJ mol?1. Estimations of the four rate constants at 50°C give, respectively, 1.2 × 104 M?3s?1, 590 M?1s?1, 2 × 109 M?2s?1, and 20 M?1 s?1. Mechanisms of these reactions involving the protonation IO2H + H+ ? IO2H and an explanation of the decrease of the last two rate constants when the temperature increases, are proposed. © 2008 Wiley Periodicals, Inc. Int J Chem Kinet 40: 647–652, 2008 相似文献
6.
Determinations of the [Ti(IV)]/[Ti(III) ratio in solutions of titanium(IV) chloride equilibrated with H2(g), at 25°C in 3 M (Na)Cl ionic medium, have indicated the predominance of the Ti(OH)22+ species in the concentration ranges 0.5 ? [H+] ? 2 M and 1.5 x 10?3 ? [Ti(IV)] ? 0.05 M. From the equilibrium data the reduction potential has been evaluated Ti(OH)22+ + 2 H+ + e ? Ti3+ + 2H2O, EoH = (7.7 ± 0.6) x 10?3 V. The acidification reactions of Ti(OH)22+ were also studied in 12 M(Li)Cl medium at 25°C by measuring the redox potential of the Ti(IV)/Ti(III) couple as a function of [H+]. The potentiometric data in the acidity range 0.3 ? [H+] ? 12 M have been explained by assuming Ti4+ + e ? Ti3+, Eo = 0.202 ± 0.002 V Ti4+ + H2O ? TiOH3+ + H+, log Ka1 = 0.3 ± 0.01 Ti4+ + 2H2O ? Ti(OH)22+ + 2H+, log Ka1Ka2 = 1.38 ± 0.05. 相似文献
7.
The equilibrium constant for the reaction CH2(COOH)2 + I3? ? CHI(COOH)2 + 2I? + H+, measured spectrophotometrically at 25°C and ionic strength 1.00M (NaClO4), is (2.79 ± 0.48) × 10?4M2. Stopped-flow kinetic measurements at 25°C and ionic strength 1.00M with [H+] = (2.09-95.0) × 10?3M and [I?] = (1.23-26.1) × 10?3M indicate that the rate of the forward reaction is given by (k1[I2] + k3[I3?]) [HOOCCH2COO?] + (k2[I2] + k4[I3?]) [CH(COOH)2] + k5[H+] [I3?] [CH2(COOH)2]. The values of the rate constants k1-k5 are (1.21 ± 0.31) × 102, (2.41 ± 0.15) × 101, (1.16 ± 0.33) × 101, (8.7 ± 4.5) × 10?1M?1·sec?1, and (3.20 ± 0.56) × 101M?2·sec?1, respectively. The rate of enolization of malonic acid, measured by the bromine scavenging technique, is given by ken[CH2(COOH)2], with ken = 2.0 × 10?3 + 1.0 × 10?2 [CH2(COOH)2]. An intramolecular mechanism, featuring a six-member cyclic transition state, is postulated to account for the results on the enolization of malonic acid. The reactions of the enol, enolate ion, and protonated enol with iodine and/or triodide ion are proposed to account for the various rate terms. 相似文献
8.
Benzylmalonic acid (BzMA) reacts via an enol mechanism with aqueous iodine to form rather stable iodobenzylmalonic acid. In the presence of iodate, which oxidizes iodide to di‐iodine, the reaction goes to completion. The kinetics of the reaction BzMA + I2 + IO3? + H+ have been studied and the results were interpreted with a suitable mechanism. In a mixture with acidic iodate, hydrogen peroxide, and manganous ions, BzMA can serve as a substrate for a Briggs–Rauscher‐type oscillating reaction. The behavior of this oscillator has been studied in detail. A mechanistic interpretation of the oscillations based on a new kinetic model is also given. © 2002 Wiley Periodicals, Inc. Int J Chem Kinet 34: 357–365, 2002 相似文献
9.
The iodine–sulfur (IS) thermochemical process for hydrogen production is one of the most promising approaches in using high‐temperature process heat supplied by a nuclear reactor. This process includes three reactions that form a closed cycle: the Bunsen reaction, in which iodine, water, and sulfur dioxide react to form sulfuric acid and hydriodic acid (HI); HI decomposition; and sulfuric acid decomposition. However, the side reactions between H2SO4 and HI may disturb the operation of the IS closed cycle. For optimal process conditions, the reaction kinetics between H2SO4 and HI should be examined. In this work, a preliminary kinetic study was conducted. Using the initial reaction rate method, the kinetic parameters of the reaction between sulfuric acid and HI, such as the apparent reaction orders and rate constant were determined. For I?, the apparent reaction order was approximately 1.77, whereas the orders for H+ and SO42? were 7.78 and 1.29, respectively. The apparent rate constant at 85 ± 1°C was approximately 2.949 × 10?11 min?1 (mol/L)?9.84. The H+ concentration had more significant influence on the reaction rate than those of SO42? and I?. Such basic data provide useful information for related process design and further kinetics study. 相似文献
10.
The complexation reactions of iron(III) with 2-pyridine carboxylic acia (picolinic acid) and 2,6-pyridine dicarboxylic acid (dipicolinic acid) in aqueous solutions have been studied by spectrophotometric and stopped flow techniques. Equilibrium constants were determined for the 1 : 1 complexes at temperatures between 25 and 80°C. The values obtained are: Picolinic Acid (HL): Fe3++ H2L+? FeHL3++H+(K1 = 2.8,ΔH = 2 kcal mole?1 at 25°C, μ = 2.67 M) Dipicolinic Acid (H2D): Fe3++H2D? FeD++2H+(K1K1A= 227 M, ΔH = 3.4 kcal mole?1 at 25°C,μ = 1.0 M). The rate constants for the formation of these complexes are also given. The results are used to evaluate the effects of these two acids upon the rate of dissolution of iron(III) from its oxides. 相似文献
11.
The solubility of H2S at 25°C in solvents of the composition: [H+]=H M, [Na+]=(I?H)=A M, [ClO4 ?]=I M was investigated by iodometric determination of [H2S]tot in the saturated solutions. Kp12=[H2S]tot·p H2S ?1 was calculated. The results are consistent with the equation:
$$\begin{gathered} \lg [H_2 S]_{tot} \cdot p_{H_2 S}^{ - 1} = --- 0,991_8 --- 0,059_0 [Na + ] + 0,008_1 [H + ]--- \hfill \\ ---0,000_1 [H + ]^4 . \hfill \\ \end{gathered} $$ 相似文献
12.
The kinetics of oxidation of Fe2+ by [Co(C3H2O4)3]3? in acidic solutions at 605 nm showed a simple first-order dependence in each reactant concentration. The second-order rate constant dependence on [H+] is in accordance with eqn (i) k2 = k′2 + k3[H+] (i) where k′2 and k3 have values of 73.4 ± 14.0 M ?1 s?1 and 353 ± 41 M?2 s?1, respectively, at 1.0 M ionic strength (NaClO4) and 25°C. At 310 nm the formation and decomposition of an intermediate, believed to be [FeC3H2O4]+, was observed. The increase in the rate of oxidation with increasing [H+] was interpreted in terms of a “one-ended” dissociation mechanism which facilitates chelation of Fe2+ by the carbonyl oxygens of malonate in the transition state. 相似文献
13.
Oxygen evolution reactions on SrFeO3 were investigated in alkaline and acidic solutions. It was found that the catalytic activity for the oxygen evolution reaction in the alkaline solution is high. The following reaction steps (V)+Fe+2H2O→(O)+FeOH2+2H++2e? in acidic solution and FeOH+OH?→FeO?+H2O in alkaline solution are presumed to be rate-controlling in the anodic evolution of oxygen on SrFeO3 electrode, where (V) denotes oxygen vacancy on the electrode surface. The reaction mechanism and the catalytic property are discussed in connection to the band structure of the oxide. 相似文献
14.
o-Phthalic acid is proposed as a standard substance for buffer solutions of known hydrogen ion concentration (I ? 0.2 M KCl, p[H+] = 3.0–5.4, 25°C). Its crystallinity, purity and slightly wide buffer range afford advantages over acetic acid. Empirical relationships between measured pH (pHm) and calculated [H+] were derived for sequences of buffer solutions at several ionic strengths: pHm - Mp[H+] + C. These calibration lines were parallel and of unit slope as required by theory. A table of p[H+] values for o-phthalic acid buffer solutions at I = 0.1 M (KCl) is presented and the method of calculation of p[H+] values for a buffer series generated by additions of potassium hydroxide is outlined. 相似文献
15.
The kinetics of the bromate oxidation of tris(1,10-phenanthroline)iron(II) (Fe(phen)32+) and aquoiron(II) (Fe2+ (aq)) have been studied in aqueous sulfuric acid solutions at μ = 1.0M and with Fe(II) complexes in great excess. The rate laws for both reactions generally can be described as -d [Fe(II)]/6dt = d[Br?]/dt = k[Fe(II)] [BrO?3] for [H+]0 = 0.428–1.00M. For [BrO?3]0 = 1.00 × 10?4M. [Fe2+]0 = (0.724–1.45)x 10?2 M, and [H+]0 = 1.00M, k = 3.34 ± 0.37 M?1s?1 at 25°. For [BrO?3]0 = (1.00–1.50) × 10?4M, [Fe2+]0 = 7.24 × 10?3M ([phen]0 = 0.0353M), and [H+]0 = 1.00M, k = (4.40 ± 0.16) × 10?2 M?1s?1 at 25°. Kinetic results suggest that the BrO?3-Fe2+ reaction proceeds by an inner-sphere mechanism while the BrO?3-Fe(phen)32+ reaction by a dissociative mechanism. The implication of these results for the bromate-gallic acid and other bromate oscillators is also presented. 相似文献
16.
The electrochemical behavior of iodine in an ambient temperature molten salt system, aluminum chloride-N-(1-butyl)pyridinium chloride (BuPyCl), have been studied in basic (excess BuPyCl) and neutral (1.0:1.0 AlCl3: BuPyCI mole ratio) melt compositions. Acid-base interactions of iodine in different oxidation states with the ionic solvent are observed. High stability of triiodide ion in neutral butylpyridinium tetrachloroaluminate indicates relatively weak intermolecular interactions in this solvent. In basic solutions polyhalogen equilibria involving iodine in different oxidation states and chloride ions are established. In iodine and tetraethylammonium triiodide solutions a mixture of ICI2?, I2Cl?, I3? and I? ions forms. The formation constants of I2Cl? and I3? and the equilibrium constant for I2Cl? disproportionation are estimated. 相似文献
17.
In aqueous H2SO4, Ce(IV) ion oxidizes rapidly Arnold's base((p-Me2NC6H4)2CH2, Ar2CH2) to the protonated species of Michler's hydrol((p-Me2NC6H4)2CHOH, Ar2CHOH) and Michler's hydrol blue((p-Me2NC6H4)2CH+, Ar2CH+). With Ar2CH2 in excess, the rate law of the Ce(IV)-Ar2CH2 reaction in 0.100 M H2SO4 is expressed -d[Ce(IV)]/dt = kapp[Ar2CH2]0[Ce(IV)] with kapp = 199 ± 8M?1s?1 at25°C. When the consumption of Ce(IV) ion is nearly complete, the characteristic blue color of Ar2CH+ ion starts to appear; later it fades relatively slowly. The electron transfer of this reaction takes place on the nitrogen atom rather than on the methylene carbon atom. The dissociation of the binuclear complex [Ce(III)ArCHAr-Ce(III)] is responsible for the appearance of the Ar2CH+ dye whereas the protonation reaction causes the dye to fade. In highly acidic solution, the rate law of the protonation reaction of Michler's hydrol blue is -d[Ar2CH+]/dt = kobs[Ar2CH+] where Kobs = ((ac + 1)[H*] + bc[H+]2)/(a + b[H+]) (in HClO4) and kobs= ((ac + 1 + e[HSO4?])[H+] + bc[H+]2 + d[HSO4?] + q[HSO4?]2/[H+])/(a + b[H+] + f[HSO4?] + g[HSO4?]/[H+]) (in H2SO4), and at 25°C and μ = 0.1 M, a = 0.0870 M s, b = 0.655 s, c = 0.202 M?1s?1, d = 0.110, e = 0.0070 M?1, f = 0.156 s, g = 0.156 s, and q = 0.124. In highly basic solution, the rate law of the hydroxylation reaction of Michler's hydrol blue is -d[Ar2CH+]/dt = kOH[OH?]0[Ar2CH+] with kOH = 174 ± 1 M?1s?1 at 25°C and μ = 0.1 M. The protonation reaction of Michler's hydrol blue takes place predominantly via hydrolysis whereas its hydroxylation occurs predominantly via the path of direct OH attack. 相似文献
18.
From solubility measurements on hydrogen sulfide in aqueous solutions of the composition [H+] = HM, [Na+] = (I — H) M = A M, [Cl?] = I M at 25°C, the molar and molal activity coefficients of H2S have been calculated. The activity coefficients of H2S in the electrolyte mixtures have been found to be additive functions of the activity coefficients in the individual electrolyte solutions at the same ionic strength. This result is predicted by the internal pressure theory of salt effects on non-electrolyte activity coefficients, provided that the volume change upon mixing two electrolyte solutions of the same ionic strength is zero. 相似文献
19.
The reaction of Cl?, Br?, I?, Co(CN)63? and NCS? with meso-tetrakis (p-trimethylammoniumphenyl)porphinatodiaquorhodate(III), [RhTAPP(H2O)2]5+, has been studied at 15, 25 and 35°C in 0.1 M [H+] with μ = 1.00 M (NaNO3). The value of the acidity constant, Kal, at 25°C is 4.39 × 10?9 M. The reactions are first order in anion concentration up to 0.9 M. The values of the stability constants, K1, and the second order rate constants, k1, for the reaction with Cl?, Br?, I?, Co(CN)63? and NCS? are respectively 0.23 M?1 and 2.5 × 10?3 M?1 s?1, 1.1 M?1 and 6.92 × 10?3 M?1 s?1, 40.0 M?1 and 17.0 × 10?3 M?1 s?1, 550 M?1 and 20.0 × 10?3 M?1 s?1, 3400 M?1 and 20.9 × 10?3 M?1 s?1. The porphine greatly labilizes the Rh(III). There has been about a 500-fold increase in the rate constant for substitution compared to that of [Rh(NH3)5H2O]3+. The substitution rates are however about the same as for [Rh(TPPS)(H2O)2]3?, indicating that the overall charge on the complex plays only a minor role. The kinetic results indicate that dissociative activation is occurring in these reactions. 相似文献
20.
Thermometric titrations of lanthanum perchlorate, titanium (III)-chloride, uranium (IV)-sulfate, and uranyl sulfate with EDTA solutions were carried out by using a Keithley nanovoltmeter with a rhodium-platinum thermocouple at 25°±0.01°. The formation of LaY?, TiY?, U(IV)Y and UO2HY? ions was confirmed. The heat of reaction for the system, Ti(III)+H2Y2? = TiY?+2H+, was given by δH1 = 1.933-1.422×10 m +2.056×104m (in cal) and the limiting value was evaluated to be δH°1 = 1.9 kcal mol?1 at 25°C. 相似文献