Nitrate aqueous equilibria

In particular, when a solvent extraction system of uranyl nitrate, aqueous nitric acid, and tri-n-butyl phosphate (TBP) in a hydrocarbon diluent was irradiated with a CO2 laser, a change was observed in the equilibrium distribution of uranyl nitrate between the phases (14). When the solution was irradiated at 944 cm 1, close to the uranyl asymmetric stretching frequency, the effect was observed. When a nonresonant frequency was used or the energy was absorbed in the solvent, no effect was observed. Little heating could be expected, and, in any case, heating effects should have been in the opposite direction from that observed. 

The state of aqueous solutions of nitric acid In strongly acidic solutions water is a weaker base than its behaviour in dilute solutions would predict, for it is almost unprotonated in concentrated nitric acid, and only partially protonated in concentrated sulphuric acid. The addition of water to nitric acid affects the equilibrium leading to the formation of the nitronium and nitrate ions ( 2.2.1). The intensity of the peak in the Raman spectrum associated with the nitronium ion decreases with the progressive addition of water, and the peak is absent from the spectrum of solutions containing more than about 5% of water a similar effect has been observed in the infra-red spectrum. ... 

Ridd - has reinterpreted the results concerning the anticatalysis of the first-order nitration of nitrobenzene in pure and in partly aqueous nitric acid brought about by the addition of dinitrogen tetroxide. In these media this solute is almost fully ionised to nitrosonium ion and nitrate ion. The latter is responsible for the anticatalysis, because it reduces the concentration of nitronium ion formed in the following equilibrium ... 

Since the first-order rate constant for nitration is proportional to y, the equilibrium concentration of nitronium ion, the above equations show the way in which the rate constant will vary with x, the stoichiometric concentration of dinitrogen tetroxide, in the two media. An adequate fit between theory and experiment was thus obtained. A significant feature of this analysis is that the weak anticatalysis in pure nitric acid, and the substantially stronger anticatalysis in partly aqueous nitric acid, do not require separate interpretations, as have been given for the similar observations concerning nitration in organic solvents. 

Electrodes of the first kind have only limited application to titration in non-aqueous media a well-known example is the use of a silver electrode in the determination of sulphides and/or mercaptans in petroleum products by titration in methanol-benzene (1 1) with methanolic silver nitrate as titrant. As an indicator electrode of the second kind the antimony pH electrode (or antimony/antimony trioxide electrode) may be mentioned its standard potential value depends on proton solvation in the titration medium chosen cf., the equilibrium reaction on p. 46). 

Pertechnetate in neutral and alkaline media can be extracted into solutions of tetra-alkylammonium iodides in benzene or chloroform. With tetra-n-heptylammo-nium iodide (7.5 x 10 M) in benzene distribution coefficients up to 18 can be obtained . A solution of fV-benzoyl-iV-phenylhydroxylamine (10 M) in chloroform can be used to extract pertechnetate from perchloric acid solution with a distribution coefficient of more than 200, if the concentration of HCIO is higher than 6 M The distribution of TcO between solutions of trilauryl-ammonium nitrate in o-xylene and aqueous solutions of nitrate has been measured. In 1 M (H, Li) NOj and 0.015 M trilaurylammonium nitrate the overall equilibrium constant has been found to be log K = 2.20 at 25 °C. The experiments support an ion exchange reaction . Pertechnetate can also be extracted with rhodamine-B hydrochloride into organic solvents. The extraction coefficient of Tc (VII) between nitrobenzene containing 0.005 %of rhodamine-B hydrochloride and aqueous alcoholic " Tc solution containing 0.0025 % of the hydrochloride, amounts to more than 5x10 at pH 4.7 . 

Bis [(trifluoromethyl)thio] acetaldehyde (83a) has been prepared from an enam-ine precursor (84), although refluxing in aqueous ethanolic HCl is required to effect this reaction.The aldehyde is less stable than its enol tautomer (83b), and many reactions typical of aldehydes fail. For example, addition of aqueous silver nitrate immediately yields the silver salt of (83b), rather than giving precipitation of (elemental) silver. The (trifluoromethyl)thio substituent has pseudohalogenic character and, together with the hydroxy group, stabilizes the alkene tautomer in the manner of a push-pull alkene. The enol-aldehyde equilibrium mixture in acetonitrile shows an apparent of 2.6 when titrated with aqueous hydroxide. 

A kinetic study of nitrous acid-catalyzed nitration of naphthalene with an excess of nitric acid in aqueous mixture of sulfuric and acetic acids (Leis et al. 1988) shows a transition from first-order to second-order kinetics with respect to naphthalene. (At this acidity, the rate of reaction through the nitronium ion is too slow to be significant the amount of nitrous acid is sufficient to make one-electron oxidation of naphthalene as the main reaction path.) The reaction that initially had the first-order in respect to naphthalene becomes the second-order reaction. The electron transfer from naphthalene to NO+ has an equilibrium (reversible) character. In excess of the substrate, the equilibrium shifts to the right. A cause of the shift is the stabilization of cation-radical by uncharged naphthalene. The stabilized cation-radical dimer (NaphH)2 is just involved in nitration ... 

At water vapor concentrations above the deliquescence point, the equilibrium is that between the reactant gases and aqueous ammonium nitrate. As treated in detail by Mozurkewich, the equilibrium constant, K 4-54, then depends on the solution concentrations or activities ... 

Experimentally, sodium nitrate was added to an aqueous solution of molecular ring 16, and the solution was heated at 100 °C. Then, the equilibrium between 16 and its catenated dimer 15 is pushed by the polar media toward the catenane. After self-assembling in a high yield, catenane 15 was isolated as a CIO4 salt in a high yield (Eq. 3). It was confirmed that catenane 15 thus obtained did not dissociate into two rings in aqueous solution because its framework had been locked. 

PbF2 (c). Guntz1 found 2.2 for the heat of reaction of aqueous lead nitrate with aqueous hydrofluoric acid. This gives for PbF2 (c), Qf=159.5. From equilibrium data, Jellinek and Rudat1 calculated Qf=156. 

PbS (c). Berthelot14 found the heats of reaction of aqueous hydrogen sulfide to be 13.34 and 11.34 with aqueous lead nitrate and aqueous lead acetate, respectively whence, for PbS (c), Qf=24.8 and 20.6, respectively. Thomsen15 found Q=30.98 for the reaction between aqueous lead nitrate and aqueous sodium sulfide whence, for PbS(c), Qf=22.5. From equilibrium data, Jellinek and Zakowski1 deduced Qf=20.0 Watanabe2 obtained 22.85 Jellinek and Deubel,1 26.2. 

Ag2C03 (c). Berthelot s12 data on the reaction of aqueous silver nitrate with aqueous potassium carbonate yield Q/=117 for the freshly precipitated silver carbonate, and Qf —121 for the crystallized material. The equilibrium data of Centnerszwer and Krustinson1 on the reaction, Ag2C03 (c) =Ag20 (c) +C02 (g), yield, for Ag2C03 (c), Q/= 119.9. 

Method. Hie metal chelates are prepared by extracting the metal ion from aqueous solution with 20-, 20- and 10-ml volumes of chloroform after addition of an appropriate amount of a solution of DDTC [0.22S g of the sodium salt in 75 ml of water and 25 ml of an ammonia-ammonium nitrate (1 1) buffer, 3 M in total ammonia]. The exact volumes which are used depend on whether the metal is uni-, bi- or tri-valent. The combined chloroform extracts are diluted to at least 50 ml for chromatography. The metal chelates are separated on plates of silica gel G or N which have been activated at 110 °C for 1 h. The Rp values of a number of DDTC metal chelates in a variety of solvent systems are listed in Table 4.31. The dried plates are sprayed with a solution consisting of 1 10 4A/ Pd(II), 7.0 10 5Af calcein and 0.02Mphosphate buffer [dihydrogen phosphate-hydrogen phosphate (1 1)]. This solution must be allowed to stand for 12 h in order to ensure that equilibrium is attained. For quantitative work with low concentrations, the solution of DDTC should be washed with chloroform before use. This removes fluorescent impurities which may cause interference in the chromatography. 

Suresh et al. investigated the extraction of uranium and thorium by TsBP and TiBP (isomers of TBP with branched carbon chain) as an alternative choice for TBP (47). Higher homologues of TBP, for example, THP and TEHP, were reported to have higher extraction ability with reduced tendency toward third-phase formation (50, 51). The esters with bulkier substituents in place of the butyl group were proposed to be of practical value for the process applications in uranium and thorium separation (54). The LOC of thorium in equilibrium with aqueous nitric acid-thorium nitrate was reported to decrease in the order THP > TAP > TBP. Pathak et al. showed that TEHP can be a better choice for U/Th separation compared to TBP and TsBP (55). 

As outlined above (p. 3), a reaction can be subject to microscopic diffusion control only if one of the reactive intermediates is formed from an inactive precursor in the reaction mixture. There are two sets of conditions which have provided evidence for microscopic diffusion control in nitration. One concerns solutions of nitric acid in aqueous mineral acids or organic solvents for, in most of these solutions, the stoicheiometric nitric acid is mainly present as the molecular species in equilibrium with a very small concentration of nitronium ions. A reaction between a substrate and a nitronium ion from this equilibrium concentration can, in principle, be subject to microscopic diffusion control. The other set of conditions is when the substrate is mainly present as the protonated form SH+ but when reaction occurs through a very small concentration of the neutral base S. A reaction between the neutral base and a nitronium ion can then, in principle, be subject to microscopic diffusion control even if the nitronium ions are the bulk component of the HN03/N0 equilibrium. In considering the evidence for microscopic diffusion control it is convenient to consider separately the reactions of those species involved in prototopic equilibria. 

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