Nitric acid complexes

Distribution and Diffusion Coefficients of Nitric Acid Complexes of TBP as a Function of TBP Concentration... 

The value of 0.145 for the equilibrium constant of the nitric acid complex is an average value derived from the equilibrium data of Moore [M2], Alcock et al. [Al], and Gruverman [G7j. The value of 0.0032 for the zirconium equilibrium is the average value derived from the equilibrium data in Table 4.4. The value of 0.00032 for the hafnium equilibrium is derived from the separation factor of 10 measured for zirconium-hafnium mixtures by Hure and Saint James [H4]. Distribution coefficients are then given by the following equations. 

Concentrated nitric acid has been used to obtain nickel(III) complexes like [Ni L(N03)2]CI04 starting from [NPL](C104)2 (L = cyclam, Megcyclam) in aqueous solutions. Oxidation of the parent nickel(II) with nitric acid complexes also afforded the nickel(IV) species [Ni L] ( 104)2 where L is the amine-imine-oxime ligand (415) or... 

It should be noted that, after completing this work, preliminary data have been obtained showing the necessity of redetermination of permissible process parameters for ion-exchange purification in the case when the resin phase is saturated with nitric acid complexes of actinides. In particular, due to increase of nitrate ion concentration, the temperature of possible thermal explosion development decreases sharply. Therefore, it is urgent that additional studies be performed that take this factor into account. 

In Fig. 6.7 the UV-Vis spectra of the dioxouranium(VI) species extracted by the TBP/[C4mim][NTf2] (30 %, v/v) are shown at different residence times and at equilibrium (3 h). As the residence time is increased, the absorption spectra approach the equilibrium limit. There is a slight shift of the most intense peaks compared to the absorption spectra of the dioxouranium(VI) species in the nitric acid solutions (see Fig. 6.2). A shift to longer wavelengths in the absorption peaks of dioxouranium(VI) ions in ionic liquids with TBP and CMPO was also reported by Visser and Rogers (2003) and was attributed to the chemical environment in the ionic liquid which favours the formation of a dioxouranium(Vl) nitric acid complex and thus causes red-shift. 

Adamantyl derivatives which have been examined include the stable nitric acid complex of (85)." The HNO3 molecule is hydrogen-bonded to... 

Other spectral interferences include argide complexes (XAr), oxides (XO), hydrides (XH), doubly charged species (X +, interfering at half the isotopic mass of X) and complex polyatomic molecules such as the nitric acid complex H2 " N 03 on A number of methods... 

It is dissolved by aqua regia (a mixture of concentrated hydrochloric and nitric acids). The product here is chlorauricil 11) acid, HAUCI4 in the complex chloraurate ion [AuClJ gold is in oxidation state + 3, auric gold. ... 

Hence mercury is a poor reducing agent it is unlikely to be attacked by acids unless these have oxidising properties (for example nitric acid), or unless the acid anion has the power to form complexes with one or both mercury cations or Hg]", so altering the... 

E values. Nitric acid attacks mercury, oxidising it to Hg (aq) when the acid is concentrated and in excess, and to Hgf (aq) when mercury is in excess and the acid dilute. Hydriodic acid Hl(aq) attacks mercury, because mercury(Il) readily forms iodo-complexes (see below, p. 438). 

The simpler nitrop>arafIins (nitromethane, nitroethane, 1- and 2-nitroproj)ane) are now cheap commercial products. They are obtained by the vapour phase nitration of the hydrocarbons a gaseous mixture of two mols of hydrocarbon and 1 mol of nitric acid vapour is passed through a narrow reaction tube at 420-476°. Thus with methane at 476° a 13 per cent, conversion into nitro methane is obtained ethane at 420° gives a 9 1 mixture of nitroethane (b.p. 114°) and nitromethane (b.p. 102°) propane at 420° afifords a 21 per cent, yield of a complex mixture of 1- (b.p. 130-6°) and 2-nitropropane (b.p. 120°), nitroethane and nitromethane, which are separated by fractional distillation. 

Nitration in aqueous solutions of nitric acid Added water retards nitration in concentrated nitric acid without disturbing the kinetic order of the reaction. The rate of nitration of nitrobenzene was depressed sixfold by the addition of 5 % of water, (c. 3 2 mol 1 ), but because of the complexity of the equilibria involving water, which exist in these media, no simple relationship could be found between the concentration of water and its effect on the rate. 

Although no chemical reaction occurs, measurements of the freezing point and infra-red spectra show that nitric acid forms i i molecular complexes with acetic acid , ether and dioxan. In contrast, the infrared spectrum of nitric acid in chloroform and carbon tetrachloride - is very similar to that of nitric acid vapour, showing that in these cases a close association with the solvent does not occur. 

Here we have the formation of the activated complex from five molecules of nitric acid, previously free, with a high negative entropy change. The concentration of molecular aggregates needed might increase with a fall in temperature in agreement with the characteristics of the reaction already described. It should be noticed that nitration in nitromethane shows the more common type of temperature-dependence (fig. 3.1). 

The kinetics of nitration in acetic anhydride are complicated. In addition to the initial reaction between nitric acid and the solvent, subsequent reactions occur which lead ultimately to the formation of tetranitromethane furthermore, the observation that acetoxylation accompanies the nitration of the homologues of benzene adds to this complexity. 

It is probable that the nitration of anthracene with nitric acid in 7-5 % aqueous sulpholan proceeds through the rapid formation of a complex. ... 

In a back titration, a slight excess of the metal salt solution must sometimes be added to yield the color of the metal-indicator complex. Where metal ions are easily hydrolyzed, the complexing agent is best added at a suitable, low pH and only when the metal is fully complexed is the pH adjusted upward to the value required for the back titration. In back titrations, solutions of the following metal ions are commonly employed Cu(II), Mg, Mn(II), Pb(II), Th(IV), and Zn. These solutions are usually prepared in the approximate strength desired from their nitrate salts (or the solution of the metal or its oxide or carbonate in nitric acid), and a minimum amount of acid is added to repress hydrolysis of the metal ion. The solutions are then standardized against an EDTA solution (or other chelon solution) of known strength. 

Control of NO emissions from nitric acid and nitration operations is usually achieved by NO2 reduction to N2 and water using natural gas in a catalytic decomposer (123—126) (see Exhaust control, industrial). NO from nitric acid/nitration operations is also controlled by absorption in water to regenerate nitric acid. Modeling of such absorbers and the complexities of the NO —HNO —H2O system have been discussed (127). Other novel control methods have also been investigated (128—129). Vehicular emission control is treated elsewhere (see Exhaust control, automotive). 

The Bachmann process, used in the United States and in some European countries, is a simplification of a series of complex reactions. In this process, a solution of one part hexamine in 1.65 parts acetic acid, and a solution of 1.50 parts ammonium nitrate dissolved in 2.0 parts nitric acid and 5.20 parts acetic anhydride are used. The reaction may be summarized as ... 

TBP and nitric acid also tend to form a complex with each other, but at sufftcientiy high uranyl nitrate concentrations the nitric acid is mainly displaced into the aqueous phase. 

Iron (III) chloride hexahydrate [10025-77-17, FeCl36H2 0, is a brown-yeUow to orange material that crystallizes from a solution of iron or iron salt dissolved ia hydrochloric acid that coataias an oxidant such as Cfy or nitric acid. The monoclinic crystals contain the complex salt... 

Qualitative. The classic method for the quaUtative determination of silver ia solution is precipitation as silver chloride with dilute nitric acid and chloride ion. The silver chloride can be differentiated from lead or mercurous chlorides, which also may precipitate, by the fact that lead chloride is soluble ia hot water but not ia ammonium hydroxide, whereas mercurous chloride turns black ia ammonium hydroxide. Silver chloride dissolves ia ammonium hydroxide because of the formation of soluble silver—ammonia complexes. A number of selective spot tests (24) iaclude reactions with /)-dimethy1amino-henz1idenerhodanine, ceric ammonium nitrate, or bromopyrogaHol red [16574-43-9]. Silver is detected by x-ray fluorescence and arc-emission spectrometry. Two sensitive arc-emission lines for silver occur at 328.1 and 338.3 nm. 

Thiocyanates are rather stable to air, oxidation, and dilute nitric acid. Of considerable practical importance are the reactions of thiocyanate with metal cations. Silver, mercury, lead, and cuprous thiocyanates precipitate. Many metals form complexes. The deep red complex of ferric iron with thiocyanate, [Fe(SCN)g] , is an effective iadicator for either ion. Various metal thiocyanate complexes with transition metals can be extracted iato organic solvents. 

In TBP extraction, the yeUowcake is dissolved ia nitric acid and extracted with tributyl phosphate ia a kerosene or hexane diluent. The uranyl ion forms the mixed complex U02(N02)2(TBP)2 which is extracted iato the diluent. The purified uranium is then back-extracted iato nitric acid or water, and concentrated. The uranyl nitrate solution is evaporated to uranyl nitrate hexahydrate [13520-83-7], U02(N02)2 6H20. The uranyl nitrate hexahydrate is dehydrated and denitrated duting a pyrolysis step to form uranium trioxide [1344-58-7], UO, as shown ia equation 10. The pyrolysis is most often carried out ia either a batch reactor (Fig. 2) or a fluidized-bed denitrator (Fig. 3). The UO is reduced with hydrogen to uranium dioxide [1344-57-6], UO2 (eq. 11), and converted to uranium tetrafluoride [10049-14-6], UF, with HF at elevated temperatures (eq. 12). The UF can be either reduced to uranium metal or fluotinated to uranium hexafluoride [7783-81-5], UF, for isotope enrichment. The chemistry and operating conditions of the TBP refining process, and conversion to UO, UO2, and ultimately UF have been discussed ia detail (40). 

Nitric acid oxidizes antimony forming a gelantinous precipitate of a hydrated antimony pentoxide (8). With sulfuric acid an indefinite compound of low solubihty, probably an oxysulfate, is formed. Hydrofluoric acid forms fluorides or fluocomplexes with many insoluble antimony compounds. Hydrochloric acid in the absence of air does not readily react with antimony. Antimony also forms complex ions with organic acids. 

Cadmium is rapidly oxidized by hot dilute nitric acid with the simultaneous generation of various oxides of nitrogen. Unlike the ziac ion, the cadmium ion is not markedly amphoteric, and therefore cadmium hydroxide [21041-95-2] Cd(OH)2, is virtually iasoluble ia alkaline media. However, the cadmium ion forms stable complexes with ammonia as well as with cyanide and haUde ions. The metal is not attacked by aqueous solutions of alkaU hydroxide. 

Nitrating cellulose with pure HNO is the simplest method of obtaining CN. In practice, nitration does not occur with acid concentrations below 75%. At acid concentrations <75%, an unstable compound (so called Knecht compound) is formed which has been described as a molecular complex or an oxonium salt of the nitric acid (72). HNO concentrations of 75—85% yield CN with 5—8% N, which dissolve in excess acid. CN with % N of 8—10% are formed at acid concentrations of 85—89%. Above 89%, a heterogeneous nitration occurs without apparent swelling of the cellulose fibers. CN with 13.3% N can be obtained with 100% HNO. Addition of inorganic salts to 100% HNO can raise the % N to 13.9. 

Copper(I) chloride is insoluble to slightly soluble in water. SolubiUty values between 0.001 and 0.1 g/L have been reported. Hot water hydrolyzes the material to copper(I) oxide. CuCl is insoluble in dilute sulfuric and nitric acids, but forms solutions of complex compounds with hydrochloric acid, ammonia, and alkaU haUde. Copper(I) chloride is fairly stable in air at relative humidities of less than 50%, but quickly decomposes in the presence of air and moisture. 

Ammonium Nitrate Plants - In ammonium nitrate plants, wet scrubbers can be considered for prill towers and the granulation plant. Particulate emissions of 0.5 kg/t of product for the prill tower and 0.25 kg/t of product for granulation should be the target. Similar loads for ammonia are appropriate. Other effluents that originate in a nitrogenous fertilizer complex include boiler blowdown, water treatment plant backwash, and cooling tower blowdown from the ammonia and nitric acid plants. 

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