Editing PH indicator (section)

==Application==
[[File:PH indicator paper.jpg|thumb|upright|pH measurement with indicator paper]]

pH indicators are frequently employed in titrations in analytical chemistry and biology to determine the extent of a [[chemical reaction]].<ref name=":0" /> 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. For a titration, the difference between the true endpoint and the indicated endpoint is called the indicator error.<ref name=":0" />

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. The figure on the right shows indicators with their operation range and color changes.

{| class="sortable wikitable"
! Indicator
! Low pH color
! Transition<br>low end
! Transition<br>high end
! High pH color
|-
| [[Crystal violet|Gentian violet]] (Methyl violet 10B)<ref>{{Cite journal |last1=Adams |first1=Elliot Q. |last2=Rosenstein |first2=Ludwig. |title=The Color and Ionization of Crystal-Violet |date=1914 |url=https://pubs.acs.org/doi/abs/10.1021/ja02184a014 |journal=Journal of the American Chemical Society |language=en |volume=36 |issue=7 |pages=1452–1473 |doi=10.1021/ja02184a014 |hdl=2027/uc1.b3762873 |issn=0002-7863|hdl-access=free }}</ref>
| style="background:#FFE100" | yellow
| align="center" | 0.0
| align="center" | 2.0
| style="background:#8A2BE2; color:white;" | blue-violet
|-
| [[Malachite green]] ([[first transition]])
| style="background:#FFE100" | yellow
| align="center" | 0.0
| align="center" | 2.0
| style="background:limegreen; color:white;" | green
|-
| [[Malachite green]] ([[second transition]])
| style="background:limegreen; color:white;" | green
| align="center" | 11.6
| align="center" | 14.0
| colorless
|-
| [[Thymol blue]] ([[first transition]])
| style="background:red; color:white;" | red
| align="center" | 1.2
| align="center" | 2.8
| style="background:#FFE100" | yellow
|-
| [[Thymol blue]] ([[second transition]])
| style="background:#FFE100" | yellow
| align="center" | 8.0
| align="center" | 9.6
| style="background:blue; color:white;" | blue
|-
| [[Methyl yellow]]
| style="background:red; color:white;" | red
| align="center" | 2.9
| align="center" | 4.0
| style="background:#FFE100" | yellow
|- 
| [[Methylene blue]] 
| style="background:#FFFFFF" | colorless
| align="center" | 5.0
| align="center" | 9.0
| style="background:#00008B; color:white" | dark blue
|- 
| [[Bromophenol blue]]
| style="background:#FFE100" | yellow
| align="center" | 3.0
| align="center" | 4.6
| style="background:blue; color:white" | blue
|-
| [[Congo red]]
| style="background:#8A2BE2; color:white;" | blue-violet
| align="center" | 3.0
| align="center" | 5.0
| style="background:red; color:white;" | red
|-
| [[Methyl orange]]
| style="background:red; color:white;" | red
| align="center" | 3.1
| align="center" | 4.4
| style="background:#FFE100" | yellow
|-
| [[Methyl orange|Screened methyl orange]] ([[first transition]])
| style="background:red; color:white;" | red
| align="center" | 0.0
| align="center" | 3.2
| style="background:plum; color:white;" | purple-grey
|-
| Screened methyl orange ([[second transition]])
| style="background:plum; color:white;" | purple-grey
| align="center" | 3.2
| align="center" | 4.2
| style="background:limegreen; color:white;" | green
|-
| [[Bromocresol green]]
| style="background:#FFE100" | yellow
| align="center" | 3.8
| align="center" | 5.4
| style="background:blue; color:white;" | blue
|-
| [[Methyl red]]
| style="background:red; color:white;" |red
| align="center" | 4.4
| align="center" | 6.2
| style="background:#FFE100" | yellow
|-
| [[Methyl purple]]
| style="background:purple; color:white;" | purple
| align="center" | 4.8
| align="center" | 5.4
| style="background:limegreen; color:white;" | green
|-
| [[Litmus test (chemistry)|Azolitmin (litmus)]]
| style="background:red; color:white;" | red 
| align="center" | 4.5
| align="center" | 8.3
| style="background:blue; color:white;" | blue
|-
| [[Bromocresol purple]]
| style="background:#FFE100" | yellow
| align="center" | 5.2
| align="center" | 6.8
| style="background:purple; color:white" | purple
|-
| [[Bromothymol blue]] 
| style="background:#FFE100" | yellow
| align="center" | 6.0
| align="center" | 7.6
| style="background:blue; color:white;" | blue
|-
| [[Phenol red]]
| style="background:#FFE100" | yellow
| align="center" | 6.4
| align="center" | 8.0
| style="background:red; color:white;" | red
|-
| [[Neutral red]]
| style="background:red; color:white;" | red
| align="center" | 6.8
| align="center" | 8.0
| style="background:#FFE100"| yellow
|-
| [[Naphtholphthalein]]
| style="background:#FF9999" | pale red
| align="center" | 7.3
| align="center" | 8.7
| style="background:#03A89E; color:white;" | greenish-blue
|-
| [[Cresol red]]
| style="background:#FFE100" | yellow
| align="center" | 7.2
| align="center" | 8.8
| style="background:#bb0080; color:white;" | reddish-purple
|-
| [[Cresolphthalein]]
| colorless
| align="center" | 8.2
| align="center" | 9.8
| style="background:purple; color:white;" | purple
|-
| [[Phenolphthalein]] (first transition)
| colorless
| align="center" | 8.3
| align="center" | 10.0
| style="background:fuchsia; color:white;" | purple-pink
|-
| Phenolphthalein (second transition)
| style="background:magenta; color:white;" | purple-pink
| align="center" | 12.0
| align="center" | 13.0
|colorless
|-
| [[Thymolphthalein]] 
| colorless
| align="center" | 9.3
| align="center" | 10.5
| style="background:blue; color:white;" | blue
|-
| [[Alizarine Yellow R]]
| style="background:#FFE100" | yellow
| align="center" | 10.2
| align="center" | 12.0
| style="background:red; color:white;" | red
|-
| [[Indigo carmine]]
| style="background:blue; color:white;" | blue
| align="center" | 11.4
| align="center" | 13.0
| style="background:gold" | yellow
|}

===[[Universal Indicator]]===

{| class=wikitable
!pH range
!Description
!Colour
|-
|1-3
|Strong acid
| style="background:red; color:white" |Red
|-
|3 – 6
|Weak acid
| style="background:#FFE100" |Orange/Yellow
|-
|7
|Neutral
| style="background:lime" |Green
|-
|8 – 11
|Weak alkali
| style="background:blue; color:white;" |Blue
|-
|11-14
|Strong alkali
| style="background:darkviolet; color:white" |Violet/Indigo
|}

=== Precise pH measurement ===
[[File:Bromocresol green spectrum.png|thumb|upright=1.8|Absorption spectra of [[bromocresol green]] at different stages of protonation]]
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>.
:HA {{eqm}} H<sup>+</sup> + A<sup>−</sup>
The value of the [[acid dissociation constant]], p''K''<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 [[Beer's law]] to be 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.
:<math chem>\begin{align}
A_x &= [\ce{HA}]\varepsilon^x_\ce{HA} + [\ce{A-}]\varepsilon^x_\ce{A-} \\
A_y &= [\ce{HA}]\varepsilon^y_\ce{HA} + [\ce{A-}]\varepsilon^y_\ce{A-}
\end{align}</math>

These are two equations in the two concentrations [HA] and [A<sup>−</sup>]. Once solved, the pH is obtained as
:<math chem>\mathrm{pH} = \mathrm{p}K_\mathrm{a}+ \log \frac{[\ce{A-}]}{[\ce{HA}]}</math>
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<sup>−</sup> (blue), weighted for the concentration of the two species.

When a single indicator is used, this method is limited to measurements in the pH range p''K''<sub>a</sub>&nbsp;±&nbsp;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, and the indicator itself is assumed to have a negligible effect on pH.