pH

A measure of the acidity of a solution, devised by the Danish chemist S. P. L. Sørenson in 1909.¹ It is the negative logarithm (base 10) of the hydrogen ion concentration expressed in moles per liter; the lower the number, the more acidic the solution is. By convention, pH's outside the range 0 — 14 are not used.

The concentration of an acid or base in a solution affects the solution's pH. Such concentrations are often expressed in terms of normality, which is the number of gram-equivalent weights of solute per liter of solution.

1. S. P. L Sørenson.
Enzyme Studies II. The Measurement and Meaning of Hydrogen Ion Concentration in Enzymatic Processes.
Biochemische Zeitschrift, volume 21, pages 131-200 (1909).

A lengthy excerpt is available online at www.chemteam.info/Chem-History/Sorenson-article.html

Some typical pHs

hydrochloric acid (1N) 0.1
sulfuric acid (1N) 0.3
human gastric juice 1.3—3.0
lemon juice 2.1
vinegar 2.3
orange juice 3.0
wine 3.4
sauerkraut 3.5
tomatoes 4.2
black coffee 5.0
rain (unpolluted) 5.7
human saliva 5.8—7.1
cow's milk 6.9
water 7.0
human urine 7.4
egg white 7.6—9.5
sea water 8.2
baking soda in water (0.1N) 8.4
household ammonia 8.4
lye (0.1N) 13.0
potassium hydroxide (1N) 14.0

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