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| {{lowercase|title=pH indicator}}
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| {{Refimprove|date=May 2012}}
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| {{Acids and bases}}
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| A '''pH indicator''' is a [[halochromism|halochromic]] [[chemical compound]] that is added in small amounts to a [[solution]] so that the [[pH]] ([[acid]]ity or [[Base (chemistry)|basicity]]) of the solution can be determined visually. Hence a pH indicator is a [[chemical]] detector for [[hydronium]] ions (H<sub>3</sub>O<sup>+</sup>) or hydrogen ions (H<sup>+</sup>) in the [[Acid-base reaction theories|Arrhenius model]]. Normally, the indicator causes the [[colour]] of the solution to change depending on the pH. Indicators can also show change in other physical properties; for example, [[olfactory indicators]] show change in their [[odor]].
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| At 25 °C, considered the standard temperature, the pH value of a neutral solution is 7.0. Solutions with a pH value below 7.0 are considered acidic, whereas solutions with pH value above 7.0 are basic (alkaline). As most naturally occurring organic compounds are weak protolytes, [[carboxylic acid]]s and [[amine]]s, pH indicators find many applications in biology and analytical chemistry. Moreover, pH indicators form one of the three main types of indicator compounds used in chemical analysis. For the quantitative analysis of metal cations, the use of [[complexometric indicator]]s is preferred, whereas the third compound class, the [[redox indicator]]s, are used in titrations involving a redox reaction as the basis of the analysis.
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| ==Theory==
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| In and of themselves, pH indicators are frequently weak acids or weak bases. The general reaction scheme of a pH indicator can be formulated as follows:
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| :<math>HInd + H_2O \rightleftharpoons H_3O^+ + Ind^-</math>
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| Here HInd stands for the acid form and Ind<sup>-</sup> for the conjugate base of the indicator. It is the ratio of these that determines the color of the solution and that connects the color to the pH value. For pH indicators that are weak protolytes, we can write the [[Henderson-Hasselbalch equation]] for them:
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| :<math>\textrm{pH} = \textrm{pK}_{a}+ \log \frac{[\textrm{Ind}^-]}{[\textrm{HInd}]}</math>
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| The equation, derived from the [[acidity constant]], states that when pH equals the pKa value of the indicator, both species are present in 1:1 ratio. If pH is above the pKa 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 pKa value, the converse is true.
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| Usually, the color change is not instantaneous at the pKa value, but there is a pH range where a mixture of colors is present. This pH range varies between indicators, but as a rule of thumb, it falls between the pKa 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 ten times greater than the concentration of the acid, their ratio is 10:1, and consequently the pH is pKa + 1. Conversely, if there is a tenfold excess of the acid with respect to the base, the ratio is 1:10 and the pH is pKa – 1.
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| 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.
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| ==Application==
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| [[File:PH indicator paper.jpg|thumb|150px|pH measurement with indicator paper.]]
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| pH indicators are frequently employed in [[titration]]s in [[analytical chemistry]] and [[biology]] to determine the extent of a [[chemical reaction]]. Because of the [[Subjectivity|subjective]] choice (determination) of color, pH indicators are susceptible to imprecise readings. For applications requiring precise measurement of pH, a [[pH meter]] is frequently used. Sometimes a blend of different indicators is used to achieve several smooth color changes over a wide range of pH values. These commercial indicators (e.g., [[universal indicator]] and [[Hydrion paper]]s) are used when only rough knowledge of pH is necessary.
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| Tabulated below are several common laboratory pH indicators. Indicators usually exhibit intermediate colors at pH values inside the listed transition range. For example, phenol red exhibits an orange color between pH 6.8 and pH 8.4. The transition range may shift slightly depending on the concentration of the indicator in the solution and on the temperature at which it is used.
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| {| class="sortable wikitable" border="1" cellpadding="2"
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| ! Indicator
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| ! Low pH color
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| ! Transition pH range
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| ! High pH color
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| |-
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| | [[Gentian violet]] ([[Methyl violet 10B]])
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| | bgcolor="yellow" | yellow
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| | align="center" | 0.0–2.0
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| | style="background:#8A2BE2; color:white;" | blue-violet
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| |-
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| | [[Malachite green]] (first transition)
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| | bgcolor="yellow" | yellow
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| | align="center" | 0.0–2.0
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| | style="background:green; color:white;" | green
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| |-
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| | [[Malachite green]] (second transition)
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| | style="background:green; color:white;" | green
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| | align="center" | 11.6–14
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| | colorless
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| |-
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| | [[Thymol blue]] (first transition)
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| | bgcolor="red" | red
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| | align="center" | 1.2–2.8
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| | bgcolor="yellow" | yellow
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| |-
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| | [[Thymol blue]] (second transition)
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| | bgcolor="yellow" | yellow
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| | align="center" | 8.0–9.6
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| | style="background:blue; color:white;" | blue
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| |-
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| | [[Methyl yellow]]
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| | bgcolor="red" | red
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| | align="center" | 2.9–4.0
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| | bgcolor="yellow" | yellow
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| |-
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| | [[Bromophenol blue]]
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| | bgcolor="yellow" | yellow
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| | align="center" | 3.0–4.6
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| | style="background:purple; color:white" | purple
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| |-
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| | [[Congo red]]
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| | style="background:#8A2BE2; color:white;" | blue-violet
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| | align="center" | 3.0–5.0
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| | bgcolor="red" | red
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| |-
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| | [[Methyl orange]]
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| | bgcolor="red" | red
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| | align="center" | 3.1–4.4
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| | bgcolor="yellow" | yellow
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| |-
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| | Screened [[methyl orange]] (first transition)
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| | bgcolor="red" | red
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| | align="center" | 0.0–3.2
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| | bgcolor="grey" | grey
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| |-
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| | Screened [[methyl orange]] (second transition)
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| | bgcolor="grey" | grey
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| | align="center" | 3.2–4.2
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| | bgcolor="green" | green
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| |-
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| | [[Bromocresol green]]
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| | bgcolor="yellow" | yellow
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| | align="center" | 3.8–5.4
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| | style="background:#0000BB; color:white;" | blue
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| |-
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| | [[Methyl red]]
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| | bgcolor="red" |red
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| | align="center" | 4.4–6.2
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| | bgcolor="yellow" | yellow
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| |-
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| | [[Litmus test (chemistry)|Azolitmin]]
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| | bgcolor="red" | red
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| | align="center" | 4.5–8.3
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| | style="background:blue; color:white;" | blue
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| |-
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| | [[Bromocresol purple]]
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| | bgcolor="yellow" | yellow
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| | align="center" | 5.2–6.8
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| | style="background:purple; color:white" | purple
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| |-
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| | [[Bromothymol blue]]
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| | bgcolor="yellow" | yellow
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| | align="center" | 6.0–7.6
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| | style="background:blue; color:white;" | blue
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| |-
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| | [[Phenol red]]
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| | bgcolor="yellow" | yellow
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| | align="center" | 6.4–8.0
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| | style="background:red; color:black" | red
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| |-
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| | [[Neutral red]]
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| | bgcolor="red" | red
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| | align="center" | 6.8–8.0
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| | bgcolor="yellow"| yellow
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| |-
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| | [[Naphtholphthalein]]
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| | bgcolor="#FF9999" | colorless to reddish
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| | align="center" | 7.3–8.7
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| | bgcolor="#03A89E" | greenish to blue
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| |-
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| | [[Cresol Red]]
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| | bgcolor="yellow" | yellow
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| | align="center" | 7.2–8.8
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| | bgcolor="#bb0080" | reddish-purple
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| |-
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| | [[Cresolphthalein]]
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| | colorless
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| | align="center" | 8.2–9.8
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| | bgcolor="red" | red
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| |-
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| | [[Phenolphthalein]]
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| | colorless
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| | align="center" | 8.3–10.0
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| | bgcolor="fuchsia" | fuchsia
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| |-
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| | [[Thymolphthalein]]
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| | colorless
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| | align="center" | 9.3–10.5
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| | style="background:blue; color:white;" | blue
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| |-
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| | [[Alizarine Yellow R]]
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| | bgcolor="yellow" | yellow
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| | align="center" | 10.2–12.0
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| | bgcolor="red" | red
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| |}
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| === Precise pH measurement ===
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| [[File:Bromocresol green spectrum.png|thumb|400px|Absorption spectra of [[bromocresol green]] at different stages of protonation.]]
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| An indicator may be used to obtain quite precise measurements of pH by measuring absorbance quantitatively at two or more wavelengths. The principle can be illustrated by taking the indicator to be a simple acid, HA, which dissociates into H<sup>+</sup> and A<sup>-</sup>.
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| :HA {{eqm}} H<sup>+</sup> + A<sup>-</sup>
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| The value of the [[acid dissociation constant]], pK<sub>a</sub>, must be known. The [[Molar absorptivity|molar absorbance]]s, ε<sub>HA</sub> and ε<sub>A<sup>-</sup></sub> of the two species HA and A<sup>-</sup> at wavelengths λ<sub>x</sub> and λ<sub>y</sub> must also have been determined by previous experiment. Assuming that [[Beer's law]] is obeyed, the measured absorbances A<sub>x</sub> and A<sub>y</sub> at the two wavelengths are simply the sum of the absorbances due to each species.
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| :<math>A_x = [HA]\epsilon^x_{HA} + [A^-]\epsilon^x_{A^-} </math>
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| :<math>A_y = [HA]\epsilon^y_{HA} + [A^-]\epsilon^y_{A^-} </math>
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| These are two equation in the two concentrations [HA] and [A<sup>-</sup>]. Once solved, the pH is obtained as
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| ::<math>\textrm{pH} = \textrm{pK}_{a}+ \log \frac{[\textrm{A}^-]}{[\textrm{HA}]}</math>
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| If measurements are made at more than two wavelengths the concentrations [HA] and [A<sup>-</sup>] can be calculated by [[linear least squares (mathematics)|linear least squares]]. In fact a whole spectrum may be used for this purpose. The process is illustrated for the indicator [[bromocresol green]]. The observed spectrum (green) is the sum of the spectra of HA (gold) and of A- (blue), weighted for the concentration of the two species.
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| When a single indicator is used this method is limited to measurements in the pH range pK<sub>a</sub> ± 1, but this range can be extended by using mixtures of two or more indicators. Because indicators have intense absorption spectra the indicator concentration is relatively low so that it can usually be assumed that the indicator itself has negligible effect on pH.
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| == Equivalence point ==
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| In acid-base titrations, an unfitting pH indicator may induce a color change in the indicator-containing solution before or after the actual equivalence point. As a result, different equivalence points for a solution can be concluded based on the pH indicator used. This is because the slightest color change of the indicator-containing solution suggests the equivalence point has been reached. Therefore, the most suitable pH indicator has an effective pH range, where the change in color is apparent, that encompasses the pH of the equivalence point of the solution being titrated.<ref>{{Cite book|title=Chemical Principles|year=2009|publisher=Houghton Mifflin Company|location=New York|pages=319–324|author=Steven S. Zumdahl|edition=6th}}</ref>
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| == Naturally occurring pH indicators ==
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| Many plants or plant parts contain chemicals from the naturally-colored [[anthocyanin]] family of compounds. They are red in acidic solutions and blue in basic. Anthocyanins can be extracted with water or other solvents from a multitude of colored plants or plant parts, including from leaves ([[red cabbage]]); flowers ([[Pelargonium|geranium]], [[poppy]], or [[rose]] petals); berries ([[blueberry|blueberries]], [[blackcurrant]]); and stems ([[rhubarb]]). Extracting anthocyanins from household plants, especially [[red cabbage]], to form a crude pH indicator is a popular introductory chemistry demonstration.
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| [[Litmus]], used by alchemists in the Middle Ages and still readily available, is a naturally occurring pH indicator made from a mixture of [[lichen]] species, particularly ''[[Roccella tinctoria]]''. The word ''litmus'' is literally from 'colored moss' in [[Old Norse]] (see [[Litr]]). The color changes between red in acid solutions and blue in alkalis. The term 'litmus test' has become a widely used metaphor for any test that purports to distinguish authoritatively between alternatives.
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| ''[[Hydrangea macrophylla]]'' flowers can change color depending on soil acidity. In acid soils, chemical reactions occur in the soil that make [[aluminium]] available to these plants, turning the flowers blue. In alkaline soils, these reactions cannot occur and therefore aluminium is not taken up by the plant. As a result, the flowers remain pink.
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| {| class="wikitable" border="1" cellpadding="2"
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| ! Indicator
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| ! Low pH color
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| ! High pH color
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| |-
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| | [[Hydrangea]] flowers
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| | bgcolor="blue" | blue
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| | bgcolor="fuchsia" | pink to purple
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| |-
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| | [[Anthocyanins]]
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| | bgcolor="red" | red
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| | bgcolor="blue" | blue
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| |-
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| | [[Litmus]]
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| | bgcolor="red" | red
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| | bgcolor="blue" | blue
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| |}
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| <gallery>
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| Image:Blue Hydrangea.jpg|Hydrangea in acid soil
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| Image:Hortensiapink.JPG|Hydrangea in alkaline soil
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| Image:Indicateur chou rouge.jpg|A gradient of red cabbage extract pH indicator from acidic solution on the left to basic on the right.
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| </gallery> | |
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| ==See also==
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| * [[Chromophore]]
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| * [[Universal indicator]]
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| * [[Nitrazine]]
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| * [[Fecal pH test]]
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| * [[pH meter]]
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| == References ==
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| <references /> | |
| {{commons|pH indicator}}
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| * [http://www.ph-meter.info/pH-measurements-indicators Long indicator list]
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| * {{fr icon}} {{PDFlink|[http://sbeccompany.fr/sciences/chimie/indicateurs/liste_indicateurs_pH.pdf Complete indicator list]|57.3 [[Kibibyte|KiB]]<!-- application/pdf, 58773 bytes -->}}
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| [[Category:PH indicators| ]]
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| [[Category:Equilibrium chemistry]]
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| [[Category:Titration]]
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| {{Link GA|fr}}
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