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pH is a measure of the acidity or alkalinity of a solution based upon the dissociation of water. The variability of pH can have a dramatic effect on biological or physical chemical reactions (e.g., geochemical weathering processes).

Substance Approximate pH H3O+
battery acid (sulfuric acid) 0 1M
lemon juice 2 1 x 10-2M
vinegar 2.5 3 x 10-3M
coffee 5 1 x 10-5M
distilled water 7 1 x 10-7M
borax 9 1 x 10-9M
household ammonia solution 11 1 x 10-11M
1M NaOH (sodium hydroxide, lye) 14 1 x 10-14M

The pH scale was developed by Danish chemist Søren Peter Lauritz Sørensen (1868–1939) in 1909 and is generally presented as ranging from 0–14, although there are no theoretical limits on the range of the scale (e.g., there are substances with negative pH's and with pH's greater than 14) the range is generally cited as being from pH =0 to pH =14.

A solution with a pH of less than pH =7 is acidic and a solution with a pH of greater than pH =7 is basic (alkaline). The midpoint of the scale, 7, is neutral. The lower the pH of a solution, the more acidic the solution is and the higher the pH, the more basic it is.

Mathematically, the potential hydronium ion concentration (pH) is equal to the negative logarithm of the hydronium ion concentration: pH = -log [H30+]. The square brackets indicate the concentration of, in moles per liter. [H 0+ 3 ] represents the hydronium ion—esentially a water molecule with a proton attached. Thus, [H OP+ 3 ] indicates the concentration of hydronium ions in moles per liter. Hydronium ions are important participants in chemical reactions that take place in aqueous (water, H20) solutions.

Water is a weak electrolyte. Through a process termed self-ionization, a small number of water molecules in pure water dissociate (separate) in a reversible reaction to form a positively charged H+ ion and a negatively charged OH- ion. In aqueous solution, as one water molecule dissociates, another is nearby to pick up the loose, positively charged, hydrogen proton to form a positively charged hydronium ion (H3O+ ). The water molecule that lost the hydrogen proton—but that kept the hydrogen electron—becomes a negatively charged hydroxide ion (OH-).

In dilute solutions, the product of the hydronium ion concentration and the hydroxide ion equals the ion product (Kw) or dissociation constant (Kw= 1.0×10-14 at 25°C). Calculations of pH using the ion product yield a number between 0 and 14—the standard pH scale.

In a sample of pure water, the concentration of hydronium ions is equal to 1×10-7 moles per liter (0.0000001 M). The equilibrium (balance) between hydronium and hydroxide ions that results from self-ionization of water can be disturbed if other substances that can donate protons are put into solution with water.

The pH of solutions may be measured experimentally with an electronic pH meter (highly accurate pH meters can measure to 0.001 pH units) or by using acid base indicators, chemicals that change color in solutions of different pH. A crude but common test for pH involves the use of Hydrion paper strips (litmus paper) that undergo changes similar to those found in indicator solutions. For example, red litmus paper turns blue in a basic solution.

See also Acids and bases.

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