From what has been said, it will be obvious that free ions can only exist in solution, and can neither evaporate, nor separate as solids; but that in the liquid they act much like other dissolved molecules, exerting their own osmotic pressure independently of each other or of the dissolved salt, but with the limitation that a solution must always contain at the same time equal numbers of + and - ions. As a solution is diluted, more ions are liberated; as it is concentrated, more recombine to form undissociated salt. This will be made clearer by an example. In a saturated solution of sodium chloride with solid salt present, we have dissolved salt at the solution-pressure of the crystallised salt, and Na and Cl ions at the dissociation-pressure of the saturated salt solution, and neither affect the others. If we now add hydrochloric acid, it has no effect directly on the solubility of the _salt_, but as HCl dissociates largely into H and Cl, it increases the pressure of the Cl ions, and so compels the salt to recombine till the Cl pressure is reduced to its normal amount. This increases the concentration of the undissociated salt-solution, and thus salt is precipitated or crystallises out till the solution is no longer super-saturated with respect to the salt-crystals. Most chemical reactions, and especially those between acids and bases, are really reactions of the ions. Thus NaOH in dilute solution is mostly ionised into Na and OH, while HCl is similarly ionised into H and Cl. On the other hand, water ionises only very slightly. Hence, on mixture, the H and OH combine and form water, with evolution of heat, while no actual combination occurs between the Na and Cl, so long as they remain in _dilute_ solution. For this reason, the heat of neutralisation of all strong acids and bases is the same, independent of their nature, since strong acids, bases and salts are almost completely ionised. The rapidity of action, and consequently what we call the “strength” or “avidity” of an acid or base depends on the number of its free ions in solution; very weak acids and bases are very little ionised, though their salts ionise almost completely in dilute solution. On this depends the explanation of a fact of great practical importance. Hydrochloric acid, a strong acid, is almost completely ionised in solution; acetic, a weak one, very little; while sodium acetate and sodium chloride as salts are both almost completely ionised. If we add hydrochloric acid to a solution of sodium acetate, we shall have sodium-ions, acet-ions, chlorine-ions and hydrogen-ions in the solution. As the pressure of the acet-ions and the hydrogen-ions will be greater than the dissociation-pressure of acetic acid, they will combine to form it, till the pressure is equalised, and we shall have in the solution, free acetic acid slightly ionised, the sodium- and chlorine-ions of sodium chloride, and the sodium- and acet-ions of any excess of sodium acetate left.