Conductivity measurements ionic strength determination

Yang et al. [12] determined the ionization constants of primaquine by a titrimetric method and studied its coordination ratio with vitamin C. The ionization constants of primaquine in 50% (v/v) ethanol in water determined at 25 °C in the ionic strength range of 5 x 10 to 5 x 10-2 mol/L are given. The coordination ratio of primaquine to vitamin C is determined by continuous variation and mole ratio methods based on pH and conductance measurements to be 1 1, indicating that the coordination compound formed in the solution is mainly a 1 1 compound. 

The electrical conductance of a solution is a measure of its current-carrying capacity and is therefore determined by the total ionic strength. It is a nonspecific property and for this reason direct conductance measurements are of little use unless the solution contains only the electrolyte to be determined or the concentrations of other ionic species in the solution are known. Conductometric titrations, in which the species of interest are converted to non-ionic forms by neutralization, precipitation, etc. are of more value. The equivalence point may be located graphically by plotting the change in conductance as a function of the volume of titrant added. 

A number of methods have been used for determining Kg values cation selective electrodes, pH-metric methods, conductimetry, calorimetry, temperature-jump relaxation measurements, membrane conductance measurements, nuclear magnetic resonance, optical rotatory dispersion. The results listed in Tables 7—10 have been obtained by various methods and at different ionic strengths so they may not always be strictly comparable. However, the corrections are probably small and the experimental accuracy is generally the same or very similar within a certain ligand type. 

Here, the association constant KA can be determined by measuring the conductivity of the unsaturated solutions. Corrections are made by successive approximations for the effect of electrolyte concentrations on molar conductivity A and for the effect of activity coefficient on KA. Here, k/Ar% is used as the first approximation of the ionic strength. 

As discussed above a certain buffer concentration is required to perform optimal analyses. The minimum ionic strength required determines the current and Joule heating. This effect can be measured as a deviation from Ohm s law. With organic buffers the conductivity is much smaller for a given ionic strength. Consequently organic zwitterionic buffers, or at least buffers with counterions of low mobility, should be preferred especially when long capillaries have to be used. 

Since electrolytes are present as impurities in the water and surfactants used, they were accounted for in the choice of the lowest electrolyte concentration. These electrolytes create a background ionic strength and may cause serious errors in Cei determination. Their presence is controlled by measuring the electrical conductivity of the surfactant solutions. In the study of thick films a reliable lower concentration limit of 1-1 valent electrolyte proves to be less than 10 4 mol dm 3 [e.g. 73,169,170]. 

In contrast to pH measurements, conductivity measurements are not widely used in the biochemistry laboratory, despite their undoubted importance. The reason for this is not clear, since the necessary apparatus and electrodes are no more expensive than those needed to measure pH, and the measurements are rapid and straightforward. Conductivity can be used readily to determine the ionic strength of solutions. 

The particular application of the Debye-Hiickel equation to be described here refers to the determination of the true equilibrium constant K from values of the equilibrium function K at several ionic strengths the necessary data for weak acids and bases can often be obtained from conductance measurements. If the solution of the electrolyte MA is sufficiently dilute for the limiting law to be applicable, it follows from equation (40.12), for the activity coe cient of a single ionic species, that... 

The water used throughout all experiments was deionized and purified with a Millipore Super Q system. It had a pH value of about 6 and conductivity which varied between 0.05 and 0.1 /iS cm The increase in ionic strength was effected with the salts used to buffer the system (O.IM NaBr aqueous solutions). Values of pH were determined with Tacussel electrode (France). The accuracy of the measurement was to 0.05 pH unit. When needed, the pH was adjusted by the controlled addition of O.IN HCl or O.IN NaOH solutions depending on the pH desired. All inorganic chemicals were of Analyzed Reagent grade. 

Values of/x = Ac/A may be calculated from Kohlrausch s measurements of electrical conductivity of hydrochloric acid solutions. /h and fci can be evaluated from the potentiometric measurements on hydrochloric acid solutions performed by Scatchaed. These data are very reliable since the concentration chain was so arranged as to eliminate diffusion potentials. In this way, ScATCHARD determined the mean activity coefficient V/h/ci instead of the individual ion activities. If we assume that in a potassium chloride solution/ = /ci— which is plausible when we recall that both ions have the same structure—and that fci is the same in hydrochloric acid solutions and potassium chloride solutions of the same concentration, then we can calculate/h and fci in hydrochloric acid solutions. Naturally these values are not strictly correct since the effect of the potassium ions on the activity of the chloride ions probably is different from that of the hydrogen ions at the same ionic strength. In the succeeding table are given values of /x, /h, and fci calculated by the above method. 

Experimental difficulties, theoretical uncertainties, and poor planning have so conspired together as to frustrate most attempts to determine the conductances or excess conductances of the electrons in amine solvents. One of the main problems in the laboratory has been the low chemical stability of the alkali metal solutions. Their blue colour gradually fades as the solutions decompose with the formation of hydrogen, a process catalysed by impurities and especially by the platinum electrodes of the cell itself. Pyrex vessels, it was recently discovered, cause sodium contamination, and for this reason much of the early research is now of doubtful worth. The experimental problems are exacerbated in the case of methylamine, whose volatility demands the use of low temperatures at which the metals dissolve but slowly. A further problem arises in the extrapolation of the data to infinitesimal ionic strength, for the appropriate conductance function to be applied depends upon the kind of species which the solution contains. And when, after all these hazards, the limiting conductance of an alkali metal solution has finally been obtained, it turns out as often as not that it can neither be compared with values for other metals because each experimenter has worked at a different temperature, nor with the conductances of normal salts because in the excitement their measurement has been overlooked. 

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