Perchlorates complex-forming ability

ESR spectra (Table 1). The JV-cyclohexylthiosemicarbazone, 13, complex formed the expected [Fe(13-H)2] with FeCl as the counterion [141]. However, [Fe(13) (13-H)H20]C104 was isolated from ethanol. Bulkiness of the cyclohexyl group, and the perchlorate ion s greater ability to hydrogen bond are probably both important to the stability of this cation. The iron(III) center is considered six-coordinate with a tridentate 13-H, bidentate 13, and a coordinated water molecule. 

Replacing the perchlorate anion in pentaamine cobalt (III) perchlorate complexes seems to be a logical way of eliminating the potentially ecologically problematic anion. The effect of such replacement was studied by Ilyushin et al. using dinitroguanidine which was considered suitable due to its ability to form complexes, its enthalpy of formation, and its oxidizing ability [27]. 

Ammonium acetate in acetic acid converts 55a into the imine (59). The hydrogen-bonded, resonance-stabilized form shown is consistent with its high melting point and intense color. The structure is further supported by the ability of the naphthalene analog, which is more soluble, to form a stable complex with cupric perchlorate. ... 

The ratio of the size of the metal ion and the radius of the internal cavity of the macrocyclic polyether determines the stoichiometry of these complexes. The stoichiometry of these complexes also depends on the coordinating ability of the anion associated with the lanthanide. For example, 12-crown-4 ether forms a bis complex with lanthanide perchlorate in acetonitrile while a 1 1 complex is formed when lanthanide nitrate is used in the synthesis [66]. Unusual stoichiometries of M L are observed when L = 12 crown-4 ether and M is lanthanide trifluoroacetate [67]. In the case of 18-crown-6 ligand and neodymium nitrate a 4 3 stoichiometry has been observed for M L. The composition of the complex [68] has been found to be two units of [Nd(18-crown-6)(N03)]2+ and [Nd(NCh)<--)]3. A similar situation is encountered [69] when L = 2.2.2 cryptand and one has [Eu(N03)5-H20]2- anions and [Eu(2.2.2)N03]+ cations. It is important to note that traces of moisture can lead to polynuclear macrocyclic complexes containing hydroxy lanthanide ions. Thus it is imperative that the synthesis of macrocyclic complexes be performed under anhydrous conditions. 

Ligands also differ in their ability to form stable complexes. The ligands phosphate, hydroxide, carbonate, and so forth, are potent complex formers while the perchlorate, CIO4", and nitrate, NO3", ions show very little tendency to form complexes. It is for this reason that nitrate or perchlorate salts are used as swamping electrolytes in experiments where it is desirable to have a constant ionic strength. ... 

Solid, hydrated nickel(II) salts and their aqueous solutions usually contain green [Ni(OH2)6l, the electronic absorption spectmm of which was shown in Fig. 20.21 with that of [Ni(NH3)g]. Salts of the latter are typically blue, giving violet solutimis. In aqueous solution, [Ni(NH3)g] is stable only in the presence of excess NH3 without which species such as [Ni(NH3)4(OH2)2] form. The violet chloride, bromide or perchlorate salts of [Ni(en)3] are obtained as racemates, the catimi being kinetically labile (see Section 26.2). The octahedral complexes trans-[Ni(C104-0,0 )2(NCMe)2] (21.52) and frans-[Ni(C104-< )2(py)4] illustrate the ability of perchlorate ions to act as bidentate or monodentate ligands respectively. The latter complex is discussed again later. 

In this section we focus on the thermodynamics of two fundamental reaction mechanisms that determine the ability of metal cations to form complex species cation hydrolysis and solvent exchange. The thermodynamic modeling strategy is first discussed, followed by the hydrolysis of the U(IV), U(V), and U(VI) aqua ions. Finally, the solvent exchange thermodynamics of Cm(in) in dilute, perchlorate (CIO4), chloride (CF), and bromide (Br ) solution will be discussed in which the effect counter anions on changes in the Cm(ni) primary hydration number. 

The formation of outer-sphere complexes also alters the solubilities of compounds. The solubility method is one of the few techniques which can be used to study outer-sphere complexes and their stepwise formation. For example, addition of perchlorate ion to aqueous solutions of hexamminecobalt(III) perchlorate first decreases the solubility, according to the law of mass action, but then, at higher concentrations of perchlorate, the solubility increases much more than changes in ionic strength would require.These experiments were used to demonstrate the ability of perchlorate ions to form ion pairs with the cobalt(III) cations and to determine the association constant. 

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