Transcript Analytical Chemistry I Lecture Note

Version 2012 Updated on 030212 Copyright © All rights reserved
Dong-Sun Lee, Prof., Ph.D. Chemistry, Seoul Women’s University
Chapter 9
Effect of Electrolytes on
Chemical Equilibria
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] = 104 M : Kw’= Kw = 1014
If [NaCl] = 101 M : Kw’= 1014
Effect of electrolyte concentration on
concentration-based equilibrium constants.
Limiting law :
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.
Ionic strength
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 Na2SO4 + 0.1M KCl
 = (1/2){(0.1×2)(+1)2+(0.1)(–2)2+(0.1)(+1)2+(0.1)(–1)2}
= 0.4M
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.
The effect of ionic strength on solubility of salts
Ex. Mercurous iodate in distilled water
Hg2(IO3)2 = Hg22+ + 2IO3–
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
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.
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

where [C] = molarity of species C,
cf
fugacity
AG = PGfG
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.
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
1)   f  ,
0  f  1
2) Z   f  (when )
3) 1 > 2  12
Effect of ionic strength on activity coefficient
Peter Debye (1884-1966) was
born and educated in Europe but
became Professor of Chemistry
at Cornell University in 1940.
He received the 1936 Nobel
Prize in Chemistry.
Ex. Interpolation
Unknown
y interval
x2–x1 = x
x2–x3 = known x interval
y2
y3
y2–y1 = y
Known
x interval
y
y1
y2–y3 = ?
Unknown y interval
(y2–y3)/(y2–y1) = (x2–x3)/(x2–x1)
 y3 = y2– [(y2–y1)(x2–x3)/(x2–x1)]
x1
x2
x
Ex. H+:  = 0.01
x3
f = 0.914
0.025
?
0.05
0.86
f = 0.914 –[(0.914 –0.86)(0.05–0.01) / (0.05–0.025)]
= 0.894
Solubility using activity coefficients
[Ca2+] =? 0.0125M 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) = Ca2+ + 2 F–
Initial
solid
0
0
Final
solid
x
2x
Ksp = 3.9 ×10 –11
Ksp0 =[Ca2+] fCa2+ [F–]2 fF–2 = (x)(0.485)(2x)2(0.81)2 = 3.9 ×10 –11
2+
x = [Ca ] = 3.1 ×10
–4
M
Solubility of LiF in distilled water (approximation)
Ksp  [Li+][F–] =x2 = 1.7 ×10 –3
x = [Li+] = [F–] = 0.041 M
1) assume  = 0.041.

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  = 0.049.

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  = 0.050.

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 

Fe3+ +
SCN– = Fe(SCN)2+
Pale yellow colorless
Red
K = [Fe(SCN)2+ ] / [Fe3+ ][SCN–]
The equilibrium quotient of
concentrations for the reaction
Fe3+ + SCN– = Fe(SCN)2+
decreases as potassium nitrate
(KNO3) is added to the solution.
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