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**Effect of Electrolytes on Chemical Equilibria**

Version 2012 Updated on Copyright © All rights reserved Dong-Sun Lee, Prof., Ph.D. Chemistry, Seoul Women’s University Chapter 9 Effect of Electrolytes on Chemical Equilibria

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**Concentration-based equilibrium constant:**

At low electrolyte (ex. NaCl), concentration-based equilibrium constant becomes independent of the electrolyte concentration and is equal to thermodynamic equilibrium constant. concentration-based equilibrium constant : Kw’, Ksp’, Ka’ … thermodynamic equilibrium constant : Kw, Ksp, Ka If [NaCl] = 104 M : Kw’= Kw = 1014 If [NaCl] = 101 M : Kw’= 1014 Effect of electrolyte concentration on concentration-based equilibrium constants.

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**Limiting law : Limiting value :**

A relationship that approaches a constant value as some parameter (ex. the electrolyte concentration) approaches zero. Limiting value : The constant numerical value observed at limit.

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**Ionic strength = (1/2)([Ci]Zi2)**

Ionic strength, is a measure of the average electrostatic interactions among ions in an electrolyte ; it is equal to one-half the sum of the terms obtained by multiplying the molarity of each ion by its valence squared. = (1/2)([Ci]Zi2) where Ci = molarity of the ith species, Zi= its charge Ex. 0.1M Na2SO M KCl = (1/2){(0.1×2)(+1)2+(0.1)(–2)2+(0.1)(+1)2+(0.1)(–1)2} = 0.4M

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**The Effect of Ionic Charges on Equilibria**

With ionic participants, the magnitude of the electrolyte effect increases with charge. Effect of electrolyte concentration on the solubility of some salts.

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**The effect of ionic strength on solubility of salts**

Ex. Mercurous iodate in distilled water Hg2(IO3)2 = Hg IO3– Ksp =1.3 ×10–18 Ksp = [Hg22+][IO3–]2 = x(2x)2 = 1.3 ×10–18 x = [Hg22+] =6.9 ×10–7M 0.050M KNO3 + saturated soln of mercurous iodate [Hg22+] =1.0 ×10–6M Ion dissociation is increased by increasing the ionic strength

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**The salt effect (also called the electrolyte effect)**

Influence of ions on the activities of reagents. The electrolyte effect results from the electrostatic attractive and repulsive forces that exist between the ions of an electrolyte and the ions involved in an equilibrium. These forces cause each ion from the dissociated reactant to be surrounded by a sheath of solution that contains a slight excess of electrolyte ions of opposite charge.

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**Activity and activity coefficient**

activity : a thermodynamic quantity which measures the effective concentration or intensity of a particular substance in a given chemical system. For dilute, ideal solutions the activity is directly proportional to the concentration ; for ideal gases, activity is proportional to the partial pressure of the gas. The absolute activity, A : = RT nl A where = chemical potential AC = [C]fc cf fugacity AG = PGfG where [C] = molarity of species C, fc = activity coefficient P = atm. General form of equilibrium constant : aA + bB = cC + dD K o= ACcADd / AAa ABb = [C]c[D]d fCc fDd / [A] a[B]b fAa fBb At low ionic strength, activity coefficients approach unity, and the thermodynamic equilibrium constant approaches the concentration equilibrium constant.

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**Activity coefficient of ions**

Estimated Debye-Hückel equation : log f = [ – 0.51 Z2 ] / [1+(3.3 )] where f = activity coefficient of an ion of charge ±Z and size (10–9 m)in an aqueous medium of ionic strength . = 0~0.1 M range f , 0 f 1 2) Z f (when ) 3) 1 > 2 12

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**Effect of ionic strength on activity coefficient**

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Peter Debye ( ) was born and educated in Europe but became Professor of Chemistry at Cornell University in 1940. He received the 1936 Nobel Prize in Chemistry.

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** x y Unknown y interval Known x interval x1 x2 y2 y1 x3 y3**

Ex. Interpolation x2–x1 = x x2–x3 = known x interval y2–y1 = y y2–y3 = ? Unknown y interval (y2–y3)/(y2–y1) = (x2–x3)/(x2–x1) y3 = y2– [(y2–y1)(x2–x3)/(x2–x1)] Ex. H+: = f = ? f = –[(0.914 –0.86)(0.05–0.01) / (0.05–0.025)] = 0.894

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**Solubility using activity coefficients**

[Ca2+] =? M MgSO4 soln saturated with CaF2. 0.0125M MgSO4 =(1/2)[(0.0125)(+2)2+(0.0125)(–2) 2] fCa2+ = 0.485, fF– = 0.81 CaF2 (s) = Ca F– Ksp = 3.9 ×10 –11 Initial solid Final solid x x Ksp0 =[Ca2+] fCa2+ [F–]2 fF–2 = (x)(0.485)(2x)2(0.81)2 = 3.9 ×10 –11 x = [Ca2+] = 3.1 ×10 –4 M

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**Solubility of LiF in distilled water (approximation) **

Ksp [Li+][F–] =x2 = 1.7 ×10 –3 x = [Li+] = [F–] = M 1) assume = fLi+= 0.851, fF– = 0.830 Ksp =[Li+] fLi+ [F–] fF– = (x)(0.851)(x)(0.830) = 1.7 ×10 –3 x = [Li+] = 0.049M 2) assume = fLi+= 0.837, fF– = 0.812 Ksp =[Li+] fLi+ [F–] fF– = (x)(0.837)(x)(0.812) = 1.7 ×10 –3 x = [Li+] = 0.050M 3) assume = fLi+= 0.835, fF– = 0.81 Ksp =[Li+] fLi+ [F–] fF– = (x)(0.835)(x)(0.81) = 1.7 ×10 –3 x = [Li+] = 0.050M solubility

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**Fe3+ + SCN– = Fe(SCN)2+ Pale yellow colorless Red**

K = [Fe(SCN)2+ ] / [Fe3+ ][SCN–] The equilibrium quotient of concentrations for the reaction Fe SCN– = Fe(SCN)2+ decreases as potassium nitrate (KNO3) is added to the solution.

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**Summary Concentration-based equilibrium**

constant , Ksp’ , Ka’, Kb’, Kw’ Ionic strength Activity, activity coefficient Fugacity, fugacity coefficient Thermodynamic equilibrium constant Electrolyte effect : salt effect Debye-Hückel equation Interpolation Solubility using activity coefficient Approximation

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