activity coefficient
Think of the activity coefficient as a fudge factor that bridges the gap between what a label says and what a solution actually does. Concentration tells you how much solute is there; the activity coefficient is the multiplier that turns that number into the effective concentration (the activity) a substance really exhibits once its particles interact.
Formally, activity = activity coefficient × concentration (or mole fraction). When a solution is ideal or infinitely dilute, the coefficient equals 1 and concentration tells the whole truth. As particles crowd in and attract or repel one another, the coefficient drifts below or above 1: a value under 1 means the solute acts weaker than its concentration suggests, a value over 1 means stronger. In effect, the coefficient is a one-number summary of how non-ideal the solution is.
This correction is what keeps quantitative chemistry honest in real solutions. Ionic solutions are the prime example: charged particles tug on one another strongly, so their activity coefficients fall well below 1 even at modest concentrations — a pattern the Debye-Hückel theory was built to predict. Equilibrium constants and electrode potentials need these coefficients to give right answers outside the dilute ideal.
In dilute salt water each ion behaves almost freely, so its activity coefficient is near 1. Crank up the concentration and the ions interfere with one another, dragging the coefficient down to, say, 0.7.
Activity coefficient = the multiplier that turns concentration into effective concentration.
Like activity, the coefficient is dimensionless and depends on the chosen reference: it approaches 1 as the solution approaches ideal behaviour. For ions, only a combined mean activity coefficient of cations and anions can be measured, since a single ion's value cannot be isolated experimentally.