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pH and Corrosion

pH , originally defined by Danish biochemist Søren Peter Lauritz Sørensen in 1909, is a measure of the concentration of hydrogen ions. The term pH was derived from the manner in which the hydrogen ion concentration is calculated, it is the negative logarithm of the hydrogen ion (H+) concentration:

pH definition

where log is a base-10 logarithm and aH+ is the activity (related to concentration) of hydrogen ions. According to the Compact Oxford English Dictionary, the "p" stands for the German word for "power", potenz, so pH is an abbreviation for "power of hydrogen".

A higher pH means there are fewer free hydrogen ions, and that a change of one pH unit reflects a tenfold change in the concentrations of the hydrogen ion. For example, there are 10 times as many hydrogen ions available at a pH of 7 than at a pH of 8. The pH scale ranges from 0 to 14. A pH of 7 is considered to be neutral. Substances with pH of less that 7 are acidic and substances with pH greater than 7 are considered to be basic.

Low pH acid waters clearly accelerate corrosion by providing a plentiful supply of hydrogen ions. Although even absolutely pure water contains some free hydrogen ions, free carbon dioxide in the water can multiply the hydrogen ion concentration many times. When carbon dioxide dissolves in water, it reacts with the water to form carbonic acid, a so-called weak acid, but an effective source of acidity. Even more acidity is sometimes encountered in acid mine waters, or in those contaminated with industrial wastes.

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See also: Calcium carbonate, Carbon dioxide, Chlorination, Dissolved oxygen, Langelier calculation, Langelier index, Larson-Skold index, Oddo-Tomson index, pH, Puckorius index, Ryznar index, Scaling Indices, Stiff-Davis index, Total dissolved solids, Water corrosivity