172Appendix: pH TheoryXL15, 20, 25, 50 and 60 metersSince its introduction by the Danish chemist Sorensen in 1909, pH measurement has becomeone of the most commonly used and important measurements in both laboratory and industrialsettings. pH measurement and control is vital to a wide array of endeavors including municipaland industrial wastewater treatment, and textile, pharmaceutical, food, and petroleum production.Even our very existence itself is dependent upon pH. Most organisms can exist only within anarrow pH range. In humans, for example, the pH of blood must be maintained within the pHrange of 7.3 to 7.4.In general, pH is a measure of the degree of acidity or alkalinity of a substance. It is related to theeffective acid concentration ("activity") of a solution by this defining equation:pH = -log aH 3O+with a H 3O+ representing the activity or effective concentration of the hydronium ion in solution.Analysts traditionally work with concentration units rather than activity. Thereforeneglecting activity, pH can be defined by the following equation:pH = -log [H 3O+]with [H 3O+] representing the concentration in Moles/liter of the hydronium ion in solution.The pH range includes values from 0 to 14. Values from 0 to 7 represent the acidic half of thescale. Values from 7 to 14 represent the alkaline or basic half of the scale. The pH value 7 isconsidered neutral, as it is neither acidic or alkaline.The pH scale is based on the dissociation constant of water. Water, even in its purest state,dissociates as follows producing a positively charged hydronium ion (H 3O+) and a negativelycharged hydroxyl ion (OH-):2H 2O = H 3O+ + OH–
