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Appendix:
pH Theory
XL15, 20, 25, 50 and 60 meters
Since its introduction by the Danish chemist Sorensen in 1909, pH measurement has become
one of the most commonly used and important measurements in both laboratory and industrial
settings. pH measurement and control is vital to a wide array of endeavors including municipal
and 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 a
narrow pH range. In humans, for example, the pH of blood must be maintained within the pH
range 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 the
effective acid concentration ("activity") of a solution by this defining equation:
pH = -log aH
3
O+
with a H
3
O+ representing the activity or effective concentration of the hydronium ion in solution.
Analysts traditionally work with concentration units rather than activity. Therefore
neglecting activity, pH can be defined by the following equation:
pH = -log [H
3
O+]
with [H
3
O+] 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 the
scale. Values from 7 to 14 represent the alkaline or basic half of the scale. The pH value 7 is
considered 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
3
O+) and a negatively
charged hydroxyl ion (OH-):
2H
2
O = H
3
O+ + OH–
Содержание XL 20
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