Anion nitrate

The distribution of highly extractable solutes such as and Pu between the aqueous and organic phases is strongly dependent upon the nitrate anion concentration in the aqueous phase. This salting effect permits extraction or reextraction (stripping) of the solute by controlling the nitric acid concentration in the aqueous phase. The distribution coefficient, D, of the solute is expressed as... 

Nitrate Anion Polystyrene matrix Trihiityl ammonium functional group Nad... 

There are notable differences in both structures and stabilities for binary N-O and S-N anions (Section 5.4). The most common oxo-anions of nitrogen are the nitrite [N02] and the nitrate anion [NOs] the latter has a branched chain structure 1.1. The sulfur analogue of nitrite is... 

The anions N3 and CN , being strong nucleophiles, form stable C—X bonds (C—N3 very polar C—CN unpolar) which can only be broken with a high energy use. The newly formed neutral compounds are relatively stable when faced with nucleophiles. The nitrate anion fits in between groups three and four and will be discussed in more detail later. 

Experimental observations and test calculations pointed out a special behaviour of the nitrate anion when faced with 6arbocations. Therefore a detailed investigation with the assistance of the MINDO/3 and the Huron-Claverie method was carried out122). It appeared that in addition to the ester formation and the proton transfer to the counterion, the formation of NO+ by oxygen transfer to the cation must be considered as well (see Fig. 11). 

The nitrate anion is a group of four atoms held together by chemical bonds. The entire unit bears a -1 electrical charge. 

Ammonium nitrate The ammonium cation always has +1 charge, and nitrate anions are always -1, so these ions combine in 1 1 ratio. The cation is listed before the anion NH4 NO3. Notice that we do not lump together the two nitrogen atoms because the ammonium and nitrate ions are distinct entities. 

Another way to calculate this molar mass is by working with the individual components of the formula. Each formula unit of the compound contains one iron(II) cation, two nitrate anions, and six waters of hydration ... 

If we were to conduct a second solubility experiment in which solutions of KI and NaN03 were mixed, we would find that no precipitate forms. This demonstrates that K and NO3 ions do not form a solid precipitate, so the bright yellow precipitate must be lead(II) iodide, Pbl2. As the two salt solutions mix, cations and r anions combine to produce lead(II) iodide, which precipitates from the solution. On standing, the yellow precipitate settles, leaving a colorless solution that contains potassium cations and nitrate anions. The molecular blowups in Figure depict these solutions at the molecular level. 

To convert from the provisional structure of the nitrate anion to complete the nitrogen octet, any of the three oxygen atoms can supply a pair of electrons. ... 

It is essential to realize that electrons In the nitrate anion do not flip back and forth among the three bonds, as implied by separate structures. The true character of the anion is a blend of the three, In which all three nitrogen-oxygen bonds are equivalent. The need to show several equivalent structures for such species reflects the fact that Lewis structures are approximate representations. They reveal much about how electrons are distributed in a molecule or ion, but they are imperfect instruments that cannot describe the entire story of chemical bonding, hi Chapter 10, we show how to interpret these structures from a more detailed bonding perspective. 

The major species present in an aqueous solution of nitric acid are water molecules, hydronium ions, and nitrate anions. The concentration of HNO3 molecules is negligible. 

Cyanide and thiocyanate anions in aqueous solution can be determined as cyanogen bromide after reaction with bromine [686]. The thiocyanate anion can be quantitatively determined in the presence of cyanide by adding an excess of formaldehyde solution to the sample, which converts the cyanide ion to the unreactive cyanohydrin. The detection limits for the cyanide and thiocyanate anions were less than 0.01 ppm with an electron-capture detector. Iodine in acid solution reacts with acetone to form monoiodoacetone, which can be detected at high sensitivity with an electron-capture detector [687]. The reaction is specific for iodine, iodide being determined after oxidation with iodate. The nitrate anion can be determined in aqueous solution after conversion to nitrobenzene by reaction with benzene in the presence of sulfuric acid [688,689]. The detection limit for the nitrate anion was less than 0.1 ppm. The nitrite anion can be determined after oxidation to nitrate with potassium permanganate. Nitrite can be determined directly by alkylation with an alkaline solution of pentafluorobenzyl bromide [690]. The yield of derivative was about 80t.with a detection limit of 0.46 ng in 0.1 ml of aqueous sample. Pentafluorobenzyl p-toluenesulfonate has been used to derivatize carboxylate and phenolate anions and to simultaneously derivatize bromide, iodide, cyanide, thiocyanate, nitrite, nitrate and sulfide in a two-phase system using tetrapentylammonium cWoride as a phase transfer catalyst [691]. Detection limits wer Hi the ppm range. 

It appeared that the chemical composition was close to the theoretical one. Surface area values decreased from Mg to Sr samples, which are related to the calcination temperatures necessary to eliminate the nitrate anions and organic compounds. Moreover, the SG samples presented higher surface area than the evaporated ones, although the calcination temperatures were higher (Table 1). 

Unlike methane and the other alkanes, aromatic hydrocarbons have absorptions in the UV part of the spectrum, and thus may be detected through UV spectrometry using silica fibers. This scheme is useful for "aromatic" water pollutants such as toluenes and xylenes with their absorption bands between 250 and 300 nm. Similarly, nitrate anion can be monitored (albeit with low sensitivity) in water via its UV absorption at 250 nm. 

The reaction of the tetranuclear mercuracarborand 166 with 2 or 3 equiv. of KNO3/I8-C-6 in acetone affords two different nitrate complexes namely [166-(N03)2(H20)]2 and [166-(N03)2]2 (Figures 15 and 16). In both cases, the nitrate anions are ligated to the mercury centers by Hg-O interactions ranging from 2.60 to 3.08 A. In [166-(N03)2]2, the two anions coordinate with all four mercury atoms in a face-on trihapto fashion from either side of the plane.216... 

The addition of mannitol to the catalyst increases the dispersion and hence the activity of the catalyst. The reason for this is the formation of a foam when large amounts of C02 and NOx gas are released following reaction of mannitol with the nitrate anion at temperatures around 80°C. We believe that the formation of this foam in the alumina pores is responsible for the increased dispersion observed in the mannitol-modified catalyst. 

The nitrate anion can also serve as the trapping agent for radical cations (181) (Eq. 2). 

Some soils, particularly those in the tropics, have significant anion exchange capacity. For these soils, there is an attraction between soil colloids and the simple halogen and nitrate anions. Bringing these anions into solution for analysis requires an extraction, or replacing anions, just as does the analysis of exchangeable cations. 

Co-crystallisation of cytosine or thymine with [Cu(L2m)][BF4]2 the tetraflu-oroborate was chosen owing to the fact that [BF4] anion forms, relative to nitrate anions, considerably weaker hydrogen-bonding interactions [67] led to [Cu(L2m)] [BF4]2 2cyt and [Cu(L2m)] [BF4]2-thy-H20,both of which adopt 2-D sheet architectures. The layered structure of [Cu(L2m)][BF4]2 2cyt is shown as a typical example in Fig. 37. In both structures, the sheets are linked by anion-mediated hydrogen-bonding interactions. 

Plutonium(V) exhibits little tendency to complex with inorganic anions. By contrast Plutonium(VI) complexes with chloride and nitrate anions but these complexes are much less stable than the corresponding Pu(IV) complexes. 

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