Constituent ion

The diffusion coefficients of the constituent ions in ionic liquids have most commonly been measured either by electrochemical or by NMR methods. These two methods in fact measure slightly different diffusional properties. The electrochemical methods measure the diffusion coefficient of an ion in the presence of a concentration gradient (Pick diffusion) [59], while the NMR methods measure the diffusion coefficient of an ion in the absence of any concentration gradients (self-diffusion) [60]. Fortunately, under most circumstances these two types of diffusion coefficients are roughly equivalent. 

For a 1-1 electrolyte (NaCl) a similar procedure may be adopted, but in this case not only do the constituent ions (Na and Cl") become more ideal in behaviour as the concentration decreases, but also the electrolyte becomes more completely dissociated into ions. It is not possible to determine indi-... 

The stability of complex ions varies within very wide limits. It is quantitatively expressed by means of the stability constant. The more stable the complex, the greater is the stability constant, i.e. the smaller is the tendency of the complex ion to dissociate into its constituent ions. When the complex ion is very stable, e.g. the hexacyanoferrate(II) ion [Fe(CN)6]4", the ordinary ionic reactions of the components are not shown. 

Both compounds are salts that dissolve in water to give their constituent ions, so the major species... 

There is difficulty in defining the absolute mobilities of the constituent ions in a molten salt, since it does not contain fixed particles that could serve as a coordinate reference. Experimental means for measuring external transport numbers or external mobilities are scarce, although the zone electromigration method (layer method) and the improved Hittorf method may be used. In addition, external mobilities in molten salts cannot be easily calculated, even from molecular dynamics simulation. 

Salts dissolve in water with dissociation of the constituent ions, this concept having been proposed originally by S. Arrhenius in 1887. His first idea was that all salts, including those of what would now be regarded as weak acids or bases, are completely dissociated at extreme dilution (Hall, 1985). It was eventually realized that substances such as NaCl, KCl, etc, are effectively completely dissociated at all concentrations. 

Some examples of compounds, along with their constituent ions shown in parentheses, are sulfuric acid (2 H+, SO4) sodium chloride (Na+, Cl ) sodium hydroxide (Na+, OH-) copper(II) sulfate (Cu2+, SO4) lead(II) nitrate (Pb2+, 2 NO3) and hydrochloric acid (H+, Cl). 

The hydrated cation Ca2+aq is of prime importance to the aqueous solution chemistry of calcium, and to most of its various roles in biological systems. The relation between lattice energy and hydration energies of the constituent ions determine solubilities, the size of the hydrated cation controls selectivity and the passage of ions through channels, and the work required to remove some or all of the water of hydration is relevant both to... 

The properties of apatites can be modified by partial or complete replacement of constituent ions - structural variations... 

The solubility product principle states that the solubility product expression for a slightly soluble compound is the product of the concentrations of its constituent ions, each raised to the power that corresponds to the number of ions in one formula unit of the compound. The quantity, K, is constant at constant temperature for a saturated solution of the compound, when the system is at equilibrium. The significance of the solubility product is that it can be used to calculate the concentrations of the ions in solutions for such slightly soluble compounds. 

The lattice enthalpy is defined as the standard change in enthalpy when a solid substance is converted from solid to form gaseous constituent ions. Accordingly, values of AH(iattice) are always positive. 

It is quite difficult to measure an accurate enthalpy of solution A//( olutioni with a calorimeter, but we can measure it indirectly. Consider the example of sodium chloride, NaCl. The ions in solid NaCl are held together in a tight array by strong ionic bonds. While dissolving in water, the ionic bonds holding the constituent ions of Na+ and Cl- in place break, and new bonds form between the ions and molecules of water to yield hydrated species. Most simple ions are surrounded with six water molecules, like the [Na(H20)6]+ ion (VI). Exceptions include the proton with four water molecules (see p. 235) and lanthanide ions with eight. 

A liquid junction potential E-f forms when the two half-cells of a cell contain different electrolyte solutions. The magnitude of Ej depends on the concentrations (strictly, the activities) of the constituent ions in the cell, the charges of each moving ion, and on the relative rates of ionic movement across the membrane. We record a constant value of j because equilibrium forms within a few milliseconds of the two half-cells adjoining across the membrane. 

In 1906, Matignon reported an enthalpy of solution of -21.54 kcal mol-1 (-90 kj mol ) for neodymium trichloride in ethanol (178). His ethanol may have been less than perfectly anhydrous, and the value for pure ethanol somewhat less negative than this, perhaps —80 or — 70 kJ mol 1. Certainly a value in this region is considerably less negative than his value for the enthalpy of solution of neodymium trichloride in water, —148 kj mol-1. The difference may reasonably be attributed to less favorable solvation qf the constituent ions in ethanol than in water. Ion solvation would be expected to be even less favorable in isopropanol, so it is not surprising to find an enthalpy of solution of about + 40 kJ mol-1 for neodymium trichloride in this alcohol. This estimate must be considered as only approximate, as it is derived from... 

We need to know which ions are available. To do this, it will be helpful to break apart any ionic reactants into their constituent ions ... 

In surface precipitation cations (or anions) which adsorb to the surface of a mineral may form at high surface coverage a precipitate of the cation (anion) with the constituent ions of the mineral. Fig. 6.9 shows schematically the surface precipitation of a cation M2+ to hydrous ferric oxide. This model, suggested by Farley et al. (1985), allows for a continuum between surface complex formation and bulk solution precipitation of the sorbing ion, i.e., as the cation is complexed at the surface, a new hydroxide surface is formed. In the model cations at the solid (oxide) water interface are treated as surface species, while those not in contact with the solution phase are treated as solid species forming a solid solution (see Appendix 6.2). The formation of a solid solution implies isomorphic substitution. At low sorbate cation concentrations, surface complexation is the dominant mechanism. As the sorbate concentration increases, the surface complex concentration and the mole fraction of the surface precipitate both increase until the surface sites become saturated. Surface precipitation then becomes the dominant "sorption" (= metal ion incorporation) mechanism. As bulk solution precipitation is approached, the mol fraction of the surface precipitate becomes large. 

The Adsorption Layer Concentration. It is not possible to measure the concentration of the constituent ions in the adsorption layer. But it has been suggested that c may be calculated by means of the Langmuir equation (10,18,19)... 

The product of the concentration of the constituent ions in a saturated solution of a difficultly soluble salt for any given temperature is practically a constant, each concentration being raised to a power equal to the relative number of ions supplied by one molecule of the salt upon dissociating . 

To understand the dissolution of ionic solids in water, lattice energies must be considered. The lattice enthalpy, A Hh of a crystalline ionic solid is defined as the energy released when one mole of solid is formed from its constituent ions in the gas phase. The hydration enthalpy, A Hh, of an ion is the energy released when one mole of the gas phase ion is dissolved in water. Comparison of the two values allows one to determine the enthalpy of solution, AHs, and whether an ionic solid will dissolve endothermically or exothermically. Figure 1.4 shows a comparison of AH and A//h, demonstrating that AgF dissolves exothermically. 

The basic questions in theory for this symposium concern the role of the forces among the constituent ions and solvent molecules in determining the thermodynamic properties of the solutions. Also there are qualitatively new thermodynamic features in some of the less well known regimes of composition and temperature. 

Crystals have typical constituent ion concentrations of about 10 so that the previous minimum concentration of Nd + ions corresponds to about 0.01 % or 100 parts per million (ppm). 

To learn that the solubility constants (products), of sparingly soluble salts can be obtained from a potentiometric titration the activity of one constituent ion is determined directly from the emf at the end point, and the salt stoichiometry then allows to be calculated. 

For example, the empirical relation between the activity and the molality ratio can be understood on the assumption that the chemical potential of the electrolyte is the sum of the chemical potentials of the constituent ions. That is, for HCl as the solute. 

The behavior of a few typical electrolytes is illustrated in Figure 19.13. By definition, 7+ is one at zero molality for all electrolytes. Furthermore, in every case, 7+ decreases rapidly with increasing molality at low values of m2. However, the steepness of this initial drop varies with the valence type of the electrolyte. For a given valence type, 7 + is substantially independent of the chemical nature of the constituent ions, as long as m2 is below about 0.01. At higher concentrations, curves for 7 + begin to separate widely and to exhibit marked specific ion effects. 

As in the case of Gibbs function changes, we also can divide the entropy change for a reaction [such as Equation (20.51)] into two parts and can assign one portion to each ion. As actual values of individual-ion entropies cannot be determined, we must establish some convention for apportioning the entropy among the constituent ions. 

The formation of a number of chromium(III) complexes Cr(H20)5 from their constituent ions Cr(H20)g and X ,... 

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