Chapter-3- Ionic
Interaction
3/5/2017 1 College Of Science and humanities, PSAU
Dr/ El Hassane ANOUAR
Chemistry Department, College of Sciences and Humanities, Prince
Sattam bin Abdulaziz University, P.O. Box 83, Al-Kharij 11942, Saudi
Arabia.
(The slides are summarized from:
Principles and applications of Electrochemistry (D. R. Crow)
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3.1 The nature of electrolytes
When electrolytes dissolved in a solvent:
The ions become free to move
The highly ordered lattice structure characteristic of crystals is
almost entirely destroyed.
Crystal structures have high lattice energies
Dissolution
In solvent
(e.g., water)
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3.1 The nature of electrolytes
Lattice energy (Elattice): The energy required to completely separate one
mole of a solid ionic compound into gaseous ions. It is always
endothermic.
Hessโ Law (Born-Haber cycles) is used to
calculate lattice energy (Elattice)
ฮH = -Elattice = ฮHf โ ฮHsub โ ยฝ ฮHBE โ IE โ EA
Lattice energy increases as
Q increases
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In a crystal:
Energies of a large number of component ion-pairs contribute to the
total lattice energy which is effectively the energy evolved when the
lattice is built up from free ions.
A large amount of energy is required to break down the ordered
structure and liberate free ions.
3.1 The nature of electrolytes
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3.1 The nature of electrolytes
Explanation of easy dissolution of lattice structure => Occurrence of
another process (Exothermic reactions of individual ions with the
solvent), which produces sufficient energy (Heat of solvation) to
compensate for that lost in the rupture of the lattice bonds.
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From the First Law of Thermodynamics:
The algebraic sum of the lattice
and solvation energies is the heat of solution.
3.1 The nature of electrolytes
This explains both:
Why the heats of solution are usually fairly small
Why they may be endothermic or exothermic -
depending upon whether the lattice energy or solvation
energy is the greater quantity.
Heat of solution = Elattice + Solvation energy
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Interionic and ion-solvent interactions
are so numerous and important in
solution that, no ion may be regarded
as behaving independently of others.
3.1 The nature of electrolytes
Certain dynamic properties such as ion conductances, mobilities and
transport numbers may be determined, although values for such
properties are not absolute but vary with ion environment.
In the most dilute cases, ion may be regarded as behaving independently
of others/
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3.2 Ion activity
In electrolytic solution, the properties of one ion species are affected by
the presence of other ions with which they interact electrostatically.
Thus, the concentration of a species is an unsatisfactory parameter to
use in attempting to predict its contribution to the bulk properties of a
solution.
We use the activity (a):
๐ฅ๐ mole fraction
๐๐ molar concentration
๐๐ molal concentration
๐พ๐ฅ rational activity coefficient
๐พ๐, ๐พ๐ are practical activity coefficients
๐๐ข = ๐๐ ๐๐ข = ๐๐ ๐๐ข = ๐๐ ๐๐ข
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3.2 Ion activity
Example
The chemical potential ฮผi of a species i may be expressed in the forms
๐๐ข = ๐๐ข๐๐ฑ + ๐๐ ๐ฅ๐ง ๐ฑ๐ข๐๐ฑ = ๐๐ข
๐๐ + ๐๐ ๐ฅ๐ง ๐๐ข๐๐ = ๐๐ข
๐๐ฆ + ๐๐ ๐ฅ๐ง๐ฆ๐ข๐๐ฆ
To determine properties of an electrolyte:
Mean ion activities (aยฑ) and mean ion activity coefficients ( ๐ยฑ)
Note: Both forms takes account of both types of ions characteristic of an electrolyte
๐ยฑ๐ = ๐+
๐+ ร ๐โ๐โ and ๐ยฑ
๐ = ๐+๐+ ร ๐โ
๐โ
๐+ Number of cations deriving from each 'molecule' of the electrolyte
๐โ Number of anions deriving from each 'molecule' of the electrolyte
๐ = ๐+ + ๐โ
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3.3 Ion-ion and ion-solvent interactions
In strong electrolytes, ions are not entirely free to move independently of
one another through in solution. Ions will move randomly with respect to
one another due to fairly violent thermal motion. Coulombic forces will
exert their influence to some extent with the result that each cation and
anion is surrounded on a time average by an 'ion atmosphere' containing
a relatively higher proportion of ions carrying charge of an opposite sign
to that on the central ion.
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3.3 Ion-ion and ion-solvent interactions
Application of an electric field:
Movement of ions will be very slow and subject to disruption by
the thermal motion.
Movement of the atmosphere occurs in a direction opposite to that
of the central ion => Breakdown symmetric ion atmosphere
As the ion moves in one direction through the solution => re-
formation of the ion atmosphere
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3.3 Ion-ion and ion-solvent interactions
The time-lag between the restructuring of the atmosphere and the movement of
the central ion causes the atmosphere to be asymmetrically distributed around
the central ion causing some attraction of the latter in a direction opposite to that
of its motion. This is known as the asymmetry, or relaxation effect.
central ions experience increased viscous hindrance to their motion on
account of solvated atmosphere ions which, on account of the latter's
movement in the opposite direction to the central ion, produce movement
of solvent in this opposing direction as well. This is known as the
electrophoretic effect.
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3.3 Ion-ion and ion-solvent interactions
The central ions experience increased viscous hindrance to their motion
on account of solvated atmosphere ions which, on account of the latter's
movement in the opposite direction to the central ion, produce movement
of solvent in this opposing direction as well. This is known as the
electrophoretic effect.
These interactions increase in significance with increasing
concentration of the electrolyte.
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2.4 The electrical potential in the vicinity of an ion
The electrical potential, ๐ , at some
point is the work done in bringing a
unit positive charge from infinity
(where ๐ = 0) to that point.
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The concentration of positive and negative ions (N+,N-) at the point P,
where the potential is ฯ may be found from the Boltzmann distribution
law, thus.
2.4 The electrical potential in the vicinity of an ion
๐+ = ๐+๐๐โ ๐ณ+๐๐ ๐ค๐
๐โ = ๐โ๐๐+ ๐ณโ๐๐ ๐ค๐
where k = Boltzmann constan
Ni = Number of ions of either kind per unit volume in the bulk
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2.4 The electrical potential in the vicinity of an ion
The electrical density (๐) at the point P where the potential is ๐ is the
excess positive or negative electricity per unit volume at that point.
๐ = ๐+๐ณ+๐ โ ๐โ๐ณโ๐
๐ = ๐+๐๐ณ+๐๐
โ ๐ณ+๐๐ ๐ค๐ โ ๐โ๐๐ณโ๐๐
+ ๐ณโ๐๐ ๐ค๐
N+ = N+0eโ z+ฯตฯ kT and Nโ = Nโ
0e+ zโฯตฯ kT
Where
Thus,
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2.4 The electrical potential in the vicinity of an ion
Example 1 : 1 electrolyte => z+ = zโ = 1 ๐๐๐ and N+0 = Nโ
0 = Ni Thus,
๐ = Niฯต eโฯตฯ kT โ eฯตฯ kT (๐. ๐
Assume that ๐๐/๐๐ โช 1
=> eโฯตฯ kT = 1 โ ฯตฯ kT and eฯตฯ kT = 1+ ฯตฯ kT
๐ ~ Niฯต 1 โ ฯตฯ kT โ 1+ ฯตฯ kT = Niฯต โ 2ฯตฯ kT
๐ ~ โ ๐๐ต๐๐๐๐ ๐๐ป
For electrolytes, ๐ง+, ๐งโ โ 1,
๐ ~ โ ๐๐ข๐ณ๐ข๐ ๐๐๐
๐ค๐
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Poisson equation: Electrostatic potential and charge density relation
๐๐๐
๐๐๐+๐๐๐
๐๐๐+๐๐๐
๐๐๐=๐๐ ๐
๐ซ
where D is the dielectric constant of the solvent medium
x, y, z are the coordinates of the point at which the potential is ฯ.
In terms of polar coordinates, Poisson equation becomes
๐
๐๐๐
๐๐๐๐๐๐
๐๐= โ๐๐ ๐
๐ซ where ๐ ~ โ ๐๐ข๐ณ๐ข
๐ ๐๐๐
๐ค๐
Thus,
๐
๐๐๐
๐๐๐๐๐๐
๐๐=๐๐
๐ซ ๐ต๐๐๐
๐ ๐๐๐
๐๐ป
2.4 The electrical potential in the vicinity of an ion
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2.4 The electrical potential in the vicinity of an ion
๐
๐๐๐
๐๐๐๐๐๐
๐๐=๐๐
๐ซ ๐ต๐๐๐
๐ ๐๐๐
๐๐ป = ๐๐๐
where
๐ = ๐๐
๐ซ ๐ต๐๐๐
๐ ๐๐
๐๐ป
๐/๐
Thus, ๐
๐๐๐
๐๐๐๐๐๐
๐๐= ๐๐๐
A general solution of Equation is
๐ =๐๐โ๐ค๐ซ
๐ซ+๐โฒ๐๐ค๐ซ
๐ซ
in which A, A' are integration constants.
Since as ๐ โ โ, ๐ โ 0 (A' = 0)
Therefore, by substitution of ๐ into ๐ expression, we obtain:
๐ = โ๐ ๐๐ข๐ณ๐ข
๐ ๐๐๐โ๐๐ซ
๐ค๐๐ซ= ๐๐๐
๐๐๐๐โ๐๐ซ
๐ซ
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For electro-neutrality:
The total negative charge of the atmosphere about a given positively
charged central ion = โ๐ง๐๐
The total charge of the atmosphere is determined by considering the charge
carried by a spherical shell of thickness dr and distance r from the central
ion and integrating from the closest distance that atmosphere and central
ions may approach out to infinity.
2.4 The electrical potential in the vicinity of an ion
Spherical
shell of
thickness
dr
Central
ion
๐๐๐ซ๐๐๐๐ซโ
๐
= ๐จ๐ฟ๐๐ซ ๐๐โ๐ฟ๐๐ ๐โ
๐
= โ๐ณ๐ข๐
๐ =๐ณ๐ข๐
๐
๐๐๐
(๐ + ๐๐)
Integration by parts
Hence,
๐ =๐ณ๐ข๐
๐
๐๐๐
(๐ + ๐๐) ๐โ๐๐ซ
๐ซ
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3.4 The electrical potential in the vicinity of an ion
Spherical
shell of
thickness
dr
Central
ion In general
๐ = ยฑ๐ณ๐ข๐
๐๐ยฑ๐ณ๐ข๐
๐
๐
๐ + ๐๐
When r approaches a, the distance of
closest approach:
๐ =๐ณ๐ข๐
๐๐
๐
(๐ + ๐๐)=๐ณ๐ข๐
๐๐โ๐ณ๐ข๐
๐
๐
(๐ + ๐๐)
๐ is composed of two contributions: ๐๐๐
๐ซ๐ Due to the ion itself
๐๐๐
๐ซ
๐ฟ
๐ + ๐ฟ๐ Represents the potential on the ion due to its atmosphere
๐ =๐ณ๐ข๐
๐
๐๐๐
(๐ + ๐๐) ๐โ๐๐ซ
๐ซ
๐ฟ
๐ + ๐ฟ๐=๐
๐ ๐ =
๐ + ๐ฟ๐
๐
The effective radius- that
of the ion atmosphere
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3.5 Electrical potential and thermodynamic functions: The
Debye-Hรผckel equation
For an ideal solution, the chemical potential ๐๐ of an ion i in is given by
๐๐ = ๐๐๐ + ๐๐ ๐๐๐๐
xi : Mole fraction of ion i
For non-ideal solutions Equation
๐๐ = ๐๐๐ + ๐๐ ๐๐๐๐ = ๐๐
๐ + ๐๐ ๐๐๐๐ + ๐๐ ๐๐๐ธ๐
By definition ฮผi is the change in free energy of the system which would occur if 1
g-ion of species i were added to a large quantity of it
๐๐ = ๐๐ ๐ธ๐
RT ln ฮณi may be regarded as the contribution that the ion atmosphere makes to the
total energy of the ion.
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The contribution per ion is ๐ค๐ ๐๐๐ธ๐ and may be equated to the work which must
be performed to give an ion of potential ฯi (due to its atmosphere) its charge ๐ง๐๐).
3.5 Electrical potential and thermodynamic functions: The
Debye-Hรผckel equation
The work done, dw, in charging the ion by an increment of charge, dฯต , is
๐๐ฐ = ๐๐ข๐๐
so that the work, w, required to give the ion its charge ๐ง๐๐ is
๐ค = ฯi
๐ง๐๐
0
๐๐ = โ ๐ง๐๐
๐ท
๐
1 + ๐ ๐
๐ง๐๐
0
๐๐ = โ๐ง๐2๐2๐
2๐ท 1 + ๐ ๐
Therefore,
๐ค๐ ๐ฅ๐ง ๐๐ข = โ๐ณ๐ข๐๐๐
๐๐
๐
๐ + ๐๐
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3.5 Electrical potential and thermodynamic functions: The
Debye-Hรผckel equation
๐ค๐ ๐ฅ๐ง ๐๐ข = โ๐ณ๐ข๐๐๐
๐๐
๐
๐ + ๐๐
In terms of the mean ion activity coefficient for the electrolyte the becomes
๐ฅ๐ง ๐ยฑ = โ๐๐
๐๐๐ค๐๐ณ+๐ณโ
๐
๐ + ๐๐
๐ =๐๐๐๐
๐๐ค๐ ๐๐ข๐ณ๐ข
๐
๐ ๐
๐๐ข =๐๐๐ข๐๐๐๐
๐ค = ๐/๐
๐ =๐๐
๐
๐๐
๐ค๐
๐
๐๐๐๐ ๐๐ข๐ณ๐ข
๐
๐ ๐
=๐๐
๐๐๐๐
๐๐๐
๐ค๐
๐
๐ ๐๐ข๐ณ๐ข
๐
๐ ๐
๐ค =๐๐
๐๐๐๐
๐๐๐
๐๐ค๐๐
๐ ๐
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๐ค =๐๐
๐๐๐๐
๐๐๐
๐๐ค๐๐
๐ ๐
3.5 Electrical potential and thermodynamic functions: The
Debye-Hรผckel equation
๐ =๐
๐ ๐๐ข๐ณ๐ข
๐ , ๐ is the ionic strength of the solution
๐ฅ๐ง ๐ยฑ = โ๐๐
๐๐๐ค๐๐ณ+๐ณโ
๐
๐ + ๐๐= โ
๐๐
๐๐ซ๐๐ป๐+๐โ
๐๐ ๐๐๐๐
๐๐๐ต๐ซ๐๐ป๐
๐ ๐
๐
๐ +๐๐ ๐๐๐๐
๐๐๐ต๐ซ๐๐ป๐๐ ๐
๐ ๐
Hence
โ ๐ฅ๐ง๐ยฑ =๐ณ+๐ณโ ๐ ๐
๐ + ๐๐ ๐
๐ต =8๐๐๐2
1000๐
1 2 1
๐ท๐ 1/2=50.29 ร 108
๐ท๐ 1/2
๐ด =๐2
2.303
2๐๐๐2
1000๐3
1 2 1
๐ท๐ 3/2=1.825 ร 106
๐ท๐ 3/2
Debye-Hรผckel equation
For water at 298 K,
A = 511 and
B = 3.29 ร 107
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2.6 Limiting and extended forms of the Debye-Hรผckel
equation
For very dilute solutions, the so-called 'Limiting Law' holds, viz.
The equation
In practice, activity coefficients show a
turning point at some value of ๐, after
which they progressively increase.
It is thus seen to be necessary to modify
Equation by the addition of a further term
which is an increasing function of ๐, i.e
๐ฅ๐ง ๐ยฑ = โ๐ ๐ณ+๐ณโ ๐
๐ + ๐๐ ๐ + ๐๐
โ ๐ฅ๐ง๐ยฑ = ๐ณ+๐ณโ ๐ ๐
โ ๐ฅ๐ง๐ยฑ =๐ณ+๐ณโ ๐ ๐
๐ + ๐๐ ๐โ โ๐ ๐ณ+๐ณโ ๐
๐ฅ๐ง ๐ยฑ= ๐ + ๐๐ ๐
Hรผckel equation
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2.7 Applications of the Debye-Hรผckel equation
The various forms of the equations resulting from the Debye-Hรผckel theory find
practical application in:
The determination of activity coefficients
The determination of thermodynamic data
2. 7.1 Determination of thermodynamte equilibrium constants
Consider dissociation of a 1: 1 weak electrolyte AB
AB โ A+ + Bโ
The thermodynamic dissociation constant KT is given by
๐พ๐ =A+ Bโ
๐ด๐ต
๐พA+๐พBโ
๐พ๐ด๐ต= ๐พ๐พยฑ2
๐พ๐ด๐ต
๐พ๐~๐พ๐พยฑ2
K: The concentration, or
conditional dissociation
constant
In dilute solution
ฮณAB = 1
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2.7 Applications of the Debye-Hรผckel equation
2. 7.1 Determination of thermodynamte equilibrium constants
AB โ A+ + Bโ ๐พ๐~๐พ๐พยฑ2
log๐พ๐ = log๐พ + 2 log ๐พยฑ
Taking
logarithms
โ ln ๐พยฑ = ๐ง+๐งโ ๐ด ๐
From the
limiting law
expression
log๐พ๐ = log๐พ โ 2๐ด ๐
log๐พ = log๐พ๐ + 2๐ด ๐
KT may therefore be determined from
measured values of over a range of ionic
strength values and extrapolating the K
versus ๐ plot to ๐ = 0.
This is a general technique for the determination of all types of thermodynamic
equilibrium constants, e.g., solubility, stability and acid dissociation constants.
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2.7 Applications of the Debye-Hรผckel equation
2. 7.2 Effect of ionic strength on ion reaction rates in solution
Consider the reaction equilibrium
AZA + BZB โ X ZA+ZB#โ Products
X ZA+ZB#: Critical complex
AZAand BZB : Reactant ions
For the pre-equilibrium
๐พ =X#
A B
๐พ#
๐พ๐ด๐พB
Note: omitting charges for clarity
The rate, ๐ฃ, of pre-equilibrium reaction:
๐ฃ = ๐ A B = ๐0 A B๐พ๐ด๐พB๐พ# where ๐ = ๐0
๐พ๐ด๐พB๐พ#
๐0 being the specific rate constant in infinitely di1ute solution, where
๐พ๐ด๐พB๐พ#= 1
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2.7 Applications of the Debye-Hรผckel equation
2. 7.2 Effect of ionic strength on ion reaction rates in solution
๐ = ๐0๐พ๐ด๐พB๐พ#
log k = log k0 + log ฮณA + log ฮณB โ log ฮณ#
โ ln ๐พยฑ =๐ง+๐งโ ๐ด ๐
1 + ๐ต๐ ๐
Express activity
coefficients as
Taking
logarithms
log ๐ = log ๐0 +๐ด ๐
1 + ๐ต๐ ๐ ๐ง๐ด2 + ๐ง๐ต
2 โ ๐ง๐ด + ๐ง๐ต2
= log๐0 +2๐ด๐ง๐ด๐ง๐ต ๐
1 + ๐ต๐ ๐
Or in very dilute solution,
log ๐ ~ log ๐0 + 2๐ด๐ง๐ด๐ง๐ต ๐
For water as solvent at 298 K log ๐ ~ log๐0 + 1.02๐ง๐ด๐ง๐ต ๐
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2.7 Applications of the Debye-Hรผckel equation
2. 7.2 Effect of ionic strength on ion reaction rates in solution
The above equations take -account of the salt effect observed for reactions
between ions, the slopes of graphs of log ๐ ๐0 versus ๐ being very close to
those predicted at low concentrations.
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At higher concentrations, deviations from linearity occur and these are
particularly noticeable for reactions having ๐ง๐ด๐ง๐ต = 0, e.g. for a reaction
between an ion and a neutral molecule.
According to Equation:
log ๐ ~ log ๐0 + 1.02๐ง๐ด๐ง๐ต ๐
such reactions should show no variation of rate with ionic strength and this is
indeed the case up until about ๐ = 0.1. Above this point, increasing ionic strength
does cause the rate to vary. The reason for this is the bฮผ term of the Hรผckel
equation. In this case it may be readily shown that
๐๐๐ ๐ = ๐๐๐ ๐0 + ๐๐ด + ๐๐ต โ ๐# ๐ (๐. ๐๐)
so that log ๐ ๐0 in this case becomes a linear function of ๐ rather than of ๐. This
has been experiment ally verified.
2.7 Applications of the Debye-Hรผckel equation
2. 7.2 Effect of ionic strength on ion reaction rates in solution
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2.8 Ion association
It is found that in many cases:
The experimental values of conductances do not agree with theoretical
values predicted by the Onsager equation:
๐๐ = ๐0 โ
8๐๐๐2
1000๐ท๐๐
1
2 ๐ง+๐งโ
3๐ท๐๐ ๐0๐
๐+๐น2๐ง๐
6๐๐๐
๐ง+ + ๐งโ
2 ๐ถ1
2
The mean ion activity coefficients cannot always be properly predicted by
the Debye-Hรผckel theory:
โ ln๐พยฑ =๐ง+๐งโ ๐ด ๐
1+๐ต๐ ๐
Under certain conditions, Bjerrum suggested that, oppositely charged ions of an
electrolyte can associate to form ion pairs. In some circumstances, even
association to the extent of forming triple or quadruple ions may occur.
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The most favourable situation for association:
Smaller ions with high charges
Solvents of low dielectric constant
2.8 Ion association
Association leads:
To a smaller number of particles in a system
Associated species have a lower charge than non-associated ones.
Diminish the magnitudes of properties of a solution which are dependent on
the number of solute particles and the charges carried by them.
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2.8 Ion association
Bjerrum's basic assumption :
The Debye-Hรผckel theory holds so long as the oppositely charged
ions of an electrolyte are separated by a distance q greater than a
certain minimum value given by
Note: When the ion separation is less than q => Ion pairing is
regarded as taking place.
q =zizjฯต2
2DkT
The above equation may be derived
from a consideration of the
Boltzmann distribution of i-type ions
in a thin shell of thickness dr at a
distance r from a central j-type ion.
Model for determination
of distribution of i-ions
within shells of specified
dimensions about j-ions
Relation of a
to q for ion
pairing
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2.8 Ion association
The number of i-type ions in such a
shell is given by
๐๐๐ข = ๐๐ข๐๐ฑ๐ฉ ยฑ๐ณ๐ข๐๐๐ฃ
๐ค๐
Model for determination
of distribution of i-ions
within shells of specified
dimensions about j-ions
Relation of a
to q for ion
pairing
The potential at a small distance from
the central j-ion may be assumed to
arise almost entirely from that ion and
is given by
๐๐ฃ = ยฑ๐ณ๐ฃ๐
๐๐ซ
Thus,
๐๐๐ข = ๐๐๐๐ข๐๐ฑ๐ฉ ยฑ๐ณ๐ข๐ณ๐ข๐
๐
๐๐ค๐๐ซ๐ซ๐๐๐ซ
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2.8 Ion association
One expects to find decreasing
probability of finding i-type ions per
unit volume at increasing values of r.
However, the volumes of the
concentric shells increase outwards
from a j-ion so that, in fact, the
probability passes through a minimum
at some critical distance.
Model for determination
of distribution of i-ions
within shells of specified
dimensions about j-ions
Relation of a
to q for ion
pairing
๐๐๐ข = ๐๐๐๐ข๐๐ฑ๐ฉ ยฑ๐ณ๐ข๐ณ๐ข๐
๐
๐๐ค๐๐ซ๐ซ๐๐๐ซ
Figure shows shapes of probability
curves for distribution of
(a) i-ions about j-ions
(b) j-ions about j-ions.
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2.8 Ion association
From the minimum condition
๐๐๐ข
๐๐ซ= ๐ โ ๐ช =
๐ณ๐ข๐ณ๐ฃ๐๐
๐๐๐ค๐
๐๐๐ข = ๐๐๐๐ข๐๐ฑ๐ฉ ยฑ๐ณ๐ข๐ณ๐ข๐
๐
๐๐ค๐๐ซ๐ซ๐๐๐ซ
For a 1:1 electrolyte in aqueous solution at 298 K, q has the value 0.357 nm.
Should the sum of the respective ionic radii be less than this figure then ion pair
formation will be favoured.
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It is evident that for a given electrolyte at constant temperature, lowering of
dielectric constant will encourage association.
Example:
For tetraisoamylammonium nitrate the sum of ion radii is of the order of 0.7 nm.
This gives a value of about 42 for D and implies that, for solvents of greater
dielectric constant than 42, there should be no association but rather complete
dissociation, i.e., the Debye-Hรผckel theory should hold good.
Conductance measurements have verified that, in fact, virtually all ion pairing has
ceased for D > 42.
2.8 Ion association