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2.1: Ionic Strength: Overview

JoVE Core
Analytical Chemistry

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Ionic Strength: Overview

2.1: Ionic Strength: Overview

The ionic strength of a solution is a quantitative way of expressing the total electrolyte concentration of a solution. This concept was first introduced in 1921 by two American physical chemists, Gilbert N. Lewis and Merle Randall, while describing the activity coefficient of strong electrolytes. During the calculation of ionic strength (I or μ), all the cations and anions are considered. However, the concentration (c) of an ion with a greater charge number (z) has a greater contribution to the total ionic strength because the charge of the ion is squared.

While calculating ionic strength for a salt that will produce multiple equivalents of the same ion on dissociation, we need to account for the contribution from each ion. For example, the ionic strength of 0.1 mol/L Na2SO4 can be calculated as follows:

The concentration of Na+ is 0.2 mol/L because one molecule of Na2SO4 will dissociate to give two Na+ ions in solution. The ionic strength of dilute solutions can be calculated easily. However, in a more concentrated solution, the calculation becomes more complex and less accurate, as the salts do not dissociate completely. For example, in an aqueous solution of 0.025 mol/L MgSO4, 25% to 35% of MgSO4 exists as the ion pair MgSO4(aq).

The concept of ionic strength can be further extended to strong and weak acids. Because strong acids dissociate completely in solution, their ionic strengths can be calculated in the same way as those of dissociated salts. For weak acids, the concentration of ionized species can be calculated from the ionization constant value and then used for ionic strength determination. If the acid is very weak and mostly remains non-ionized, its contribution to the total ionic strength of the solution is negligible.


Ionic Strength Electrolyte Concentration Activity Coefficient Cations Anions Charge Number Na2SO4 MgSO4 Strong Acids Weak Acids Ionization Constant

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