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By L. Pataki

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Of pure water and of a buffer solution on addition of an acid or base in the same amounts is illustrated in figure 2. The magnitude of the pH-stabilizing effect of buffer solutions, the buffer capacity, is the concentration of a strong acid or base, in gequivalent/dm 3 required to cause a unit change in ρΆ. Buffer capacit y is unity if 1 g-equivalent of an acid or base changes the ρΉ. of 1 dm 3 of buffer solution by one unit. If the acid-base ratio is too small or too 35 CHEMICAL EQUILIBRIA IN SOLUTION 14 12 / 1 10 8 6 4 X \2 \ 2 v»_ 005 vj O-IO 0-15 0-20 mol/dm : i Figure 2 C h a n g e s i n t h e pH.

Thus some ligands in a solution, to an extent t h a t depends on the hydrogen ion concentration, form complexes with the metal ion, while some are present in protonated form, unable to take part in complex formation. So the hydrogen ion concentration affects the concentration of the ligands capable of complex formation, thereby influencing the concentration of the complex, too. 0 a '{ a4 Vo ai a 2 n f\ Kf'D OCAJ 0Γ —4 +1 lg[NH,] 1 +2 Figure 4 Distribution curves of copper(II) ammines as a function of the ligand concentration; [Cu2 + ] = [NH 3 ] = 10- 2 mol/dm 3 If Cx is the overall concentration of the ligand, the material balance is: Cx = [ H X ] + [ X - ] .

Effect of complex formation Increased solubility in the presence of foreign electrolytes is not always the consequence of the decrease in the activity coefficients; it may often arise from complex formation. I n this case a sharp increase in solubility as the electrolyte concentration increases gives a certain indication of the reason for the increase. Whereas the decrease in the activity coefficients can increase the solubility by up to ten times, complex formation can produce solubilities up to 10 20 times higher.

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