A polyprotic acid is a compound which may lose more than 1 proton. Stepwise dissociation constants are each defined for the loss of a single proton. The constant for dissociation of the first proton may be denoted as K and the constants for dissociation of successive protons as K, etc. Phosphoric acid, , is an example of a polyprotic acid as it can lose three protons.
When the difference between successive pK values is about four or more, as in this example, each species may be considered as an acid in its own right; In fact salts of  may be crystallised from solution by adjustment of pH to about 5.5 and salts of  may be crystallised from solution by adjustment of pH to about 10. The species distribution diagram shows that the concentrations of the two ions are maximum at pH 5.5 and 10.
When the difference between successive pK values is less than about four there is overlap between the pH range of existence of the species in equilibrium. The smaller the difference, the more the overlap. The case of citric acid is shown at the right; solutions of citric acid are buffered over the whole range of pH 2.5 to 7.5.
According to Paulings first rule, successive pK values of a given acid increase . For oxyacids with more than one ionizable hydrogen on the same atom, the pK' values often increase by about 5 units for each proton removed, as in the example of phosphoric acid above.
It can be seen in the table above that the second proton is removed from a negatively charged species. Since the proton carries a positive charge extra work is needed to remove it, which is why pK is greater than pK. pK is greater than pK because there is further charge separation. When an exception to Pauling's rule is found, it indicates that a major change in structure is also occurring. In the case of (aq), the vanadium is octahedral, 6-coordinate, whereas vanadic acid is tetrahedral, 4-coordinate. This means that four "particles" are released with the first dissociation, but only two "particles" are released with the other dissociations, resulting in a much greater entropy contribution to the standard Gibbs free energy change for the first reaction than for the others.