ionic strength
Imagine measuring how 'electrically busy' a solution is — not just how many ions are floating in it, but how strongly they tug on one another. A solution of doubly charged ions feels far more crowded, electrically, than the same number of singly charged ones. Ionic strength is the single number that captures this total electrical congestion.
It is computed by adding up, for every ion in the solution, its concentration multiplied by the square of its charge, then taking half the total. Squaring the charge is the key move: it makes higher-charge ions count for much more, matching how much harder they pull on their neighbours. Pure water has an ionic strength near zero; seawater has a high one.
Ionic strength matters because it is the master dial controlling activity coefficients: the higher the ionic strength, the more each ion's effective concentration shrinks below its actual one. That is why analysts often deliberately swamp a sample with an inert salt to fix the ionic strength at a known, constant value, so that activity coefficients stay steady and calibration holds. The caveat: ionic strength counts every ion present, including spectator ions that take no part in the reaction you care about, so even an 'inert' background salt changes it.
A 0.1 M solution of NaCl (singly charged ions) has ionic strength 0.1, but a 0.1 M solution of MgSO4 (doubly charged ions) has ionic strength 0.4 — four times higher, even at the same molarity.
Charge squared: doubly charged ions raise ionic strength fast.
Ionic strength is a property of the whole solution, not of any one ion. Even an ion you do not care about contributes to it, which is precisely why a constant 'ionic strength adjustment buffer' (ISAB) is added in techniques like ion-selective electrode measurements to keep activity coefficients fixed.