Thermo Fisher Scientific CyberScan pH 5000 - Manual (Page 133)

Instruction Manual CyberScan pH5500/ 5000
15.4 pH Theory
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 industry settings. pH
measurement and control is vital to a wide array of endeavors
including municipal and industrial wastewater treatment, and textile,
pharmaceutical, food and pH. Most organisms can exit 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 aH
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 nor alkaline.
The pH scale is based on the dissociation constant of water. Water,
even in its purest state, dissociated as follows producing a positively
charged hydronium ion (H
3
O+) and a negatively charges hydroxyl ion
(OH
-
):
2H
2
O = H
3
O
+
+ OH
-
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