Solvents molar conductivities

Radii of anions of lithium salts and limiting molar conductivities in solvents of... 

In aqueous electrolyte solutions the molar conductivities of the electrolyte. A, and of individual ions, Xj, always increase with decreasing solute concentration [cf. Eq. (7.11) for solutions of weak electrolytes, and Eq. (7.14) for solutions of strong electrolytes]. In nonaqueous solutions even this rule fails, and in some cases maxima and minima appear in the plots of A vs. c (Eig. 8.1). This tendency becomes stronger in solvents with low permittivity. This anomalons behavior of the nonaqueous solutions can be explained in terms of the various equilibria for ionic association (ion pairs or triplets) and complex formation. It is for the same reason that concentration changes often cause a drastic change in transport numbers of individual ions, which in some cases even assume values less than zero or more than unity. 

Besides complexes of thiosemicarbazones prepared from nitrogen heterocycles, iron(III) complexes of both 2-formylthiophene thiosemicarbazone, 26, and 2-acetylthiophene thiosemicarbazone, 27, have been isolated [155]. Low spin, distorted octahedral complexes of stoichiometry [Fe(26)2A2]A (A = Cl, Br, SCN) were found to be 1 1 electrolytes in nitromethane. Low spin Fe(27)3A3 (A = Cl, Br, SCN) complexes were also isolated, but their insolubility in organic solvents did not allow molar conductivity measurements. Infrared speetra indicate coordination of both via the azomethine nitrogen and thione sulfur, but not the thiophene sulfur. The thiocyanate complexes have spectral bands at 2065, 770 and 470 cm consistent with N-bonded thiocyanato ligands, but v(FeCl) and v(FeBr) were not assigned due to the large number of bands found in the spectra of the two ligands. 

The molar conductance values of the complex Ln(DPSO)6 I3 in acetonitrile are slightly higher than those suggested for 1 1 electrolytes, due to the displacement of some coordinated iodide by the solvent (250). The conductance values observed for the complexes, however, approach more closely the values reported for 1 1 electrolytes as the ionic size of the lanthanide ion decreases. This may be due to the increasing strength of the metal-anion bond with decreasing cation size. 

Molar conductivity measurements are equally applicable to both solid and liquid electrolytes. In contrast, the measurement of current flowing through an electrochemical cell on a time scale of minutes or hours while the cell is perturbed by a constant dc potential is only of value for solid solvents (Bruce and Vincent, 1987) where convection is absent. Because of the unique aspects of dc polarisation in a solid solvent this topic is treated in some detail in this chapter. Let us begin by considering a cell of the form ... 

The break-seal on A is then crushed, the sample to be investigated diluted with the solvent in C, and the optical density of the resulting solution determined in the appropriate optical cell. The solution is then transferred to conductivity cell H, and its resistance is measured. The optical density is redetermined, and thereafter about two thirds of the solution is transferred to C. The solvent from C is distilled into the chilled ampoule G and used to dilute the residual solutions left in H. The conductivity and the optical density of this solution are determined as described previously thereafter, two thirds is again transferred to C, and the remaining one third is diluted by the above-described procedure. In this way the conductivities are determined for decreasing concentrations of living polymer, so that the molar conductivity A can be calculated as a function of [living polymer] down to about 10" M. 

I2]. The substantial solubilities of these compounds in chloroform and other less polar organic solvents are in agreement with their formulation as nonelectrolytes. In methanol at 25° C., the molar conductivities of 166 and 167 ohm-1 for [Ni-(NH2CH2CH2S-CH3)2I2] and [Ni(NH2CH2CH2S-CH2C6H5)2I2], respectively, are characteristic of di-univalent electrolytes in this solvent, indicating almost complete solvolysis of the coordinated iodide ions in this relatively polar solvent. Decomposition of these complexes was observed upon dissolving in water. Visible and near-infrared spectra results are also consistent with structure VI. 

Analytical data on the soluble products isolated from chloroform are in excellent agreement with the composition 1 Ni+2 1 monoalkylated ligand 1 I or Br. The magnetic moment of this methylated complex was found to be 1.89 Bohr magnetons per nickel (II). The molar conductivities of the methylated and benzylated complexes in methanol at 25° C. are 75.4 and 68.4 ohm-1, respectively. These values approximate those expected for uni-univalent electrolytes in this solvent. The formulation of these alkylated compounds as dimeric electrolytes (structure VII) does not appear to be totally consistent with their physical properties. One or both halide ions may be bound to the metal ion. These results lead to the easily understood generalization that terminal sulfur atoms alkylate more readily than bridged mercaptide groups. 

In the above sections, we considered electrolytes that are ionophores.10 Iono-phores, like sodium chloride, are ionic in the crystalline state and are expected to dissociate into free ions in dilute solutions. In fact, in high-permittivity solvents (er>40), ionophores dissociate almost completely into ions unless the solutions are of high concentration. When an ionophore is completely dissociated in the solution, its molar conductivity A decreases linearly with the square root of the concentration c (<10 2 M) ... 

With the decrease in permittivity, however, complete dissociation becomes difficult. Some part of the dissolved electrolyte remains undissociated and forms ion-pairs. In low-permittivity solvents, most of the ionic species exist as ion-pairs. Ion-pairs contribute neither ionic strength nor electric conductivity to the solution. Thus, we can detect the formation of ion-pairs by the decrease in molar conductivity, A. In Fig. 2.12, the logarithmic values of ion-association constants (log KA) for tetrabutylammonium picrate (Bu4NPic) and potassium chloride (KC1) are plotted against (1 /er) [38]. 

Tab. 6.5 Ionic molar conductivities (A00) in some organic solvents and LJPs between solutions in the same solvent (Ej) calculated by the Henderson equation...

When the limiting molar conductivities are to be obtained for a series of ions in a given solvent, the first step is to get the limiting molar conductivity of an ion by one of the above methods. Then, the limiting molar conductivities for other ions can be obtained sequentially by applying Kohlrausch s law of independent ionic migration (Section 5.8). 

The limiting molar conductivities of ions in various solvents are listed in Table 7.4. The following are some general points about ionic conductivities in non-aque-ous solutions ... 

The molar conductivities of alkali metal ions increase, in most solvents, with increasing crystal ionic radii (Li+[Pg.216]

In dipolar aprotic solvents, the molar conductivity of the fastest anion is larger than that of the fastest cation. However, this does not apply in protic solvents. 

Kundu and Bhattacharya14 have isolated dioxouranium complexes of benzohydroxamic acid with the compositions M[U02(C7H602N)3] [where M = Li, Na, K, Cs, Tl, N4, pyH+ (pyridinium) or agH+ (aminoguanidinium)] and M [U02(C7H602N)3]2 [where M = enH2+ (ethylene-diammonium)]. All the complexes, with the exception of the sodium compound, are insoluble in common organic solvents but are soluble in DMSO and DMF. The complexes have been characterized on the basis of electronic, IR and molar conductance data in DMF. Their fairly stable character is indicated by thermogravimetric analysis and the stability order is NH4+ < Tl+ < Cs+ < Li+ w Na+ agH+ < K+ pyH+ < enHl+. 

The protected amino acid tetrachlorophthalimidoacetic acid synthesised from tetra-chlorophthalic anhydride and glycine using DMF as the solvent was conducted on a molar scale to yield 300 g of the product after 8 min (Scheme 9.8)113. 

Most of the salts are hydrolyzed in moist air but are stable indefinitely in a dry nitrogen atmosphere. A notable exception is tetraethylammonium hexabromotantalate(V), which is quite stable in the atmosphere and is only slowly hydrolyzed by concentrated aqueous ammonia. Complete hydrolysis is more readily achieved by the addition of acetone followed by aqueous ammonia, the mixture being gently warmed on a water bath. They are generally insoluble in nonpolar or slightly polar solvents but are fairly soluble in more polar solvents such as acetonitrile (methyl cyanide) in the latter solvent the solutions show the expected molar conductivities. 

The color of the crystals is pale yellow to almost white. A pink shade shows the presence of minor impurities. The quality of the product is conveniently judged from the ultraviolet spectrum. The bands at wave numbers 35,000 and 44,100 cm.-1 with molar extinction coefficients of 73 and 333, respectively, are typical of the chloropentaammine. Spectra of pink crystals show less resolved bands. These products are, however, sufficiently pure to serve as starting materials for further reactions. The molar conductivity in aqueous solution is 253 cm.2 ohm-1 mole-1. The crystals are soluble in water but insoluble in HC1 or in common organic solvents. 

The relatively low molar conductance is associated with the high viscosity of this solvent. 

Pig. 3. Molar conductivities of Sn(CH3)3l (c = 7 x 10" mole/liter) dissolved in nitrobenzene on addition of increasing amounts of various electron pair donor (EPD) solvents. HMPA, hexamethylphosphoric amide DMSO, dimethyl sulfoxide DMF, dimethylformamide TBP, tributyl phosphate THF, tetrahydrofuran AN, acetonitrile. 

Sn(CH3)3l dissolved in nitrobenzene as a function of concentration of various EPD solvents added (35). In noncoordinating or weakly coordinating solvents, such as hexane, earbon tetrachloride, 1,2-dichloroethane, nitrobenzene, or nitromethane, Sn(CH3)3l is present in an unionized state (tetrahedral molecules). Addition of a stronger EPD solvent to this solution provokes ionization, presumably with formation of trigonal bipju amidal cations [Sn(CH3)3 (EPD)2J. Table II reveals that the molar conductivities at a given mole ratio EPD Sn(CH3)3l are (with the exception of pyridine) in accordance with the relative solvent donicities. No relationship appears to exist between conductivities and the dipole moments or the dielectric constants of the solvents. 

Ni(PEX)Br2 is intermediate in properties and behavior. In nonpolar solvents, such as dichloroethane, the compound exists as a neutral, molecular, 6-coordinate species having a normal, triplet, ground state, as inferred from molecular weight, magnetic moment, and electronic spectra. However, in nitromethane, at room temperature, the substance exhibits intermediate values for molar conductance (58 ohm at 10 M) and magnetic moment (2.65 BM at 9.44 x 10 M). Dilution experiments yield values for the molar conductance that are consistent with an equilibrium constant of 1.66 X 10 3 (T = 25°C.) for the process given in Equation 8. 

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