por The ionic force is a way of expressing the concentration of ions in a solution. This concept was introduced in 1922 by Lewis and Randall while they were working on the description of the coefficient of chemical activity.
When the concentration of ions in a solution is high, an electrostatic interaction occurs between the oppositely charged ions; that is, cations and anions attract each other very strongly, which results in the actual or effective ionic concentration being less than that calculated for a particular chemical reaction.
For this reason, the concept of chemical activity as the effective ionic concentration of a solution was introduced, the chemical activity being the product of the molarity of the solution times the coefficient of chemical activity.
This coefficient has a value close to unity (1) for dilute ionic solutions and for so-called ideal solutions. These are solutions where the intermolecular interaction between similar molecules is equal to that between dissimilar molecules.
The creation of the concept of ionic strength contributed to the explanation of the deviations from the ideal behavior, observed in real ionic solutions.
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ionic strength units
Ionic strength has units of moles/L (molarity) or moles/Kg of water (molality). The latter is recommended in non-ideal solutions, which are characterized because the volumes of their mixtures are not totally additive.
This means, for example, the following: if 0.5 liters of liquid A and 0.5 liters of liquid B are mixed, the resulting volume of this mixture will not necessarily be equal to 1 liter, but may be different.
Ionic strength is represented by the symbol I.
How to calculate ionic strength?
To calculate the ionic strength of a solution, the concentration of all the ions present in the solution, as well as their respective valences, are taken into account.
The ionic strength value is obtained by applying the following formula:
Where I, as already stated, is the ionic strength; C, corresponds to the molar or molal ionic concentration of the ions; while Z represents their respective valences (±1, ±2, ±3, etc.).
The expression that appears in the formula when calculating ionic strength (Σ) is read as summation, that is, the sum of the product of the molar concentration (C) of each ion present in the solution times its high valence (Z). squared.
As can be seen, the valence of the ion has the greatest weight in the value of the ionic strength of the solution. For example: the valence (Z) of Ca is +2, so Z2 equals 4. Meanwhile, the valence (Z) of Na is +1, and therefore Z2 equals 1.
This indicates that the contribution of the Ca2+ ion to the ionic strength value, at the same molar ionic concentration, is four times greater than that of the Na+ ion.
Importance of ionic strength
The ionic strength is a suitable measure of the ionic concentration of a solution and is the basis of the establishment of the Debye-Hückel Theory. This theory describes the ideal behavior of ionic solutions.
The ionic strength serves as the basis for calculating the activity coefficient (γi), a parameter that in turn allows the calculation of the chemical activity of an ionic compound, the chemical activity being the effective and real concentration of an ionic compound in solution.
As the ionic strength of a solution increases, the interaction between ions increases. Therefore, γi and the chemical activity of the ions decrease.
An increase in ionic strength can decrease the solubility of proteins in an aqueous medium, this property being used for protein precipitation selectively. High ionic strength ammonium sulfate solutions are used for the precipitation and purification of plasma proteins.
Examples of ionic forces
Example 1
Calculate the ionic strength of a 0.3 M potassium chloride (KCl) solution.
KCl dissociates in the following way:
KCl → K+ + Cl–
So we have two ions: the cation K+ (Z =+1) and the anion Cl– (Z=-1). We then apply the formula to calculate the ionic strength I:
I = 1/2 [C · (+1)1 + C · (-1)1]
= 1/2 [0.3 M · 11 + 0.3 M · 11]
= 0.3M
Note that the valence -1 of Cl– was taken as 1, its absolute value, since otherwise the ionic strength would be equal to 0.
Example 2
Calculate the ionic strength of a 0.5 M calcium sulfate (CaSO4) solution.
CaSO4 dissociates in the following way:
CaSO4 → Ca2+ + SO42-
We have two ions: the cation Ca2+ (Z =+2) and the anion SO42- (Z=-2). We then apply the formula to calculate the ionic strength I:
I = 1/2 [C · (+2)2 + C · (-2)2]
= 1/2 [0,5 M · 4 + 0,5 M · 4]
= 2M
Example 3
Calculate the ionic strength of a buffer with the final concentrations of 0.3 M dibasic sodium phosphate (Na2HPO4) and 0.4 M monobasic sodium phosphate (NaH2PO4).
Na2HPO4 dissociates as follows:
Na2HPO4 → 2Na+ + HPO42-
While NaH2PO4 dissociates following the following pattern:
NaH2PO4 → Na+ + H2PO4–
We proceed as in the previous exercises, this time having the anions HPO42- (Z= -2) and H2PO4– (Z= -1):
I = 1/2 {[C · 2 · (+1)1 + C · (-2)2] + [C · (+1)1 + C · (-1)1]}
= 1/2 {[0.3 M · 2 · 1 + 0.3 M · 4] + [0.4 M · 1 + 0.4 M · 1]}
= 1/2 {[0.6 M + 1.2 M] + [0.4 M + 0.4 M]}
= 1.3M
Note that the concentration of Na+ coming from Na2HPO4 is multiplied by 2, since its concentration is double. However, for the other salt, NaH2PO4, the Na+ concentration is multiplied by 1, according to the stoichiometry of its dissolution equation.
Example 4
Calculate the ionic strength of a solution of 0.15 M sodium chloride (NaCl) and 0.3 M glucose (C6H12O6).
NaCl dissociates in the following way:
NaCl → Na+ + Cl–
Glucose, however, does not dissociate into ions because it only has covalent bonds in its chemical structure. Therefore, the valence of glucose (Z) is equal to zero (0). We then calculate the ionic strength product of NaCl:
I = 1/2 [C · (+1)1 + C · (-1)1]
= 1/2 [0.15 M · 1 + 0.15 M · 1]
= 0.15M
References
Whitten, Davis, Peck & Stanley. (2008). Chemistry. (8th ed.). CENGAGE Learning.
Wikipedia. (2020). Ionic strength. Retrieved from: en.wikipedia.or
Dr. David K. Ryan. (nd). Activity & Ionic Strength Class 4 Ryan. [PDF]. Retrieved from: faculty.uml.edu
University of Michigan. (nd). A More Detailed Look at Chemical Equilibria. [PDF]. Retrieved from: umich.edu
Elsevier BV (2020). Ionic Strength. Science Direct. Retrieved from: sciencedirect.com
CD Kennedy. (1990). Ionic Strength and the Dissociation of Acids. [PDF]. Retrieved from: iubmb.onlinelibrary.wiley.com