Editing PH indicator (section)

==Theory==
In and of themselves, pH indicators are usually weak acids or weak bases. The general reaction scheme of acidic pH indicators in aqueous solutions can be formulated as:

:HInd<sub>(aq)</sub> + {{chem|H|2|O}}<sub>(l)</sub> {{eqm}} {{chem|H|3|O<sup>+</sup>}}<sub>(aq)</sub> + {{chem|Ind<sup>−</sup>}}<sub>(aq)</sub>

where, "HInd" is the acidic form and "Ind<sup>−</sup>" is the conjugate base of the indicator.

Vice versa for basic pH indicators in aqueous solutions:

:IndOH<sub>(aq)</sub> + {{chem|H|2|O}}<sub>(l)</sub> {{eqm}} {{chem|H|2|O}}<sub>(l)</sub> + {{chem|Ind<sup>+</sup>}}<sub>(aq)</sub> + {{chem|O|H<sup>−</sup>}}<sub>(aq)</sub>

where "IndOH" stands for the basic form and "Ind<sup>+</sup>" for the [[conjugate acid]] of the indicator.

The ratio of [[concentration]] of conjugate acid/base to concentration of the acidic/basic indicator determines the pH (or pOH) of the solution and connects the color to the pH (or pOH) value. For pH indicators that are weak electrolytes, the [[Henderson–Hasselbalch equation]] can be written as:

:pH = p''K''<sub>a</sub> + [[Common logarithm|log<sub>10</sub>]] {{sfrac|&thinsp;[{{chem|Ind<sup>−</sup>}}]&thinsp;|&thinsp;[HInd]&thinsp;}}

{{center|''or''}}

:pOH = p''K''<sub>b</sub> + log<sub>10</sub> {{sfrac|&thinsp;[{{chem|Ind<sup>+</sup>}}]&thinsp;|&thinsp;[IndOH]&thinsp;}}

The equations, derived from the [[Acid dissociation constant|acidity constant]] and basicity constant, states that when pH equals the p''K''<sub>a</sub> or p''K''<sub>b</sub> value of the indicator, both species are present in a 1:1 ratio. If pH is above the p''K''<sub>a</sub> or p''K''<sub>b</sub> value, the concentration of the conjugate base is greater than the concentration of the acid, and the color associated with the conjugate base dominates. If pH is below the p''K''<sub>a</sub> or p''K''<sub>b</sub> value, the converse is true.

Usually, the color change is not instantaneous at the p''K''<sub>a</sub> or p''K''<sub>b</sub> value, but a pH range exists where a mixture of colors is present. This pH range varies between indicators, but as a rule of thumb, it falls between the p''K''<sub>a</sub> or p''K''<sub>b</sub> value plus or minus one. This assumes that solutions retain their color as long as at least 10% of the other species persists. For example, if the concentration of the conjugate base is 10 times greater than the concentration of the acid, their ratio is 10:1, and consequently the pH is p''K''<sub>a</sub>&nbsp;+&nbsp;1 or p''K''<sub>b</sub>&nbsp;+&nbsp;1. Conversely, if a 10-fold excess of the acid occurs with respect to the base, the ratio is 1:10 and the pH is p''K''<sub>a</sub>&nbsp;−&nbsp;1 or p''K''<sub>b</sub>&nbsp;−&nbsp;1.

For optimal accuracy, the color difference between the two species should be as clear as possible, and the narrower the pH range of the color change the better. In some indicators, such as [[phenolphthalein]], one of the species is colorless, whereas in other indicators, such as [[methyl red]], both species confer a color. While pH indicators work efficiently at their designated pH range, they are usually destroyed at the extreme ends of the pH scale due to undesired side reactions.