Ionic strength (I) is a standard physical-chemistry quantity that measures how crowded and how highly charged the ions in a solution are. It's the key input for estimating activity coefficients via the Debye–Hückel equation, for predicting solubility shifts (the salting-in/salting-out effect), and for characterizing buffers, seawater, and physiological fluids.

How the Ionic Strength Calculator works

The calculator implements the standard definition I = ½Σcᵢzᵢ², summing over every ion you enter. Each ion's contribution is ½cᵢzᵢ² in mol/L; the ½ factor is conventional so that each pairwise ion interaction is counted once rather than twice. Because the charge term is squared, a doubly charged ion (z = 2) contributes four times as much per mole as a singly charged ion at the same concentration — which is why 0.1 M MgCl₂ (I = 0.3 mol/L) has three times the ionic strength of 0.1 M NaCl (I = 0.1 mol/L) even though both are 0.1 molar.

Inputs and what they mean

Each ion needs a concentration in mol/L and a signed charge number. Concentration is always non-negative; charge carries a sign (positive for cations, negative for anions), but since the formula squares it, the sign itself doesn't change the contribution — only the magnitude does. On the From a Salt tab, enter the salt's own molar concentration and the calculator applies the correct stoichiometric multiplier for each ion the salt releases (for example, CaCl₂ → 1 mol Ca²⁺ + 2 mol Cl⁻ per mol of salt dissolved).

Limits and edge cases

This calculator assumes complete dissociation, which is an excellent approximation for strong electrolytes (common salts, strong acids, strong bases) but overstates ionic strength for weak or partially dissociated species, where only a fraction of the dissolved formula units actually split into free ions. It also treats every listed concentration as already known and correct — it does not account for ion pairing, complexation, or activity effects at high ionic strength, where the simple additive formula and downstream models like the Debye–Hückel limiting law (accurate below roughly I = 0.01 mol/L) begin to break down. For concentrated or complex solutions, specialized activity-coefficient models (Davies equation, Pitzer equations) are more appropriate than the raw ionic-strength value alone.