View source for PH indicator

{{Short description|Chemical added to show pH of a solution}}{{lowercase|title=pH indicator}}
{{Acids and bases}}

[[File:Acid-base-indicators.png|thumb|pH indicators: a graphic view]]
A '''pH indicator''' is a [[halochromism|halochromic]] [[chemical compound]] added in small amounts to a [[Solution (chemistry)|solution]] so the [[pH]] ([[acid]]ity or [[Base (chemistry)|basicity]]) of the solution can be determined visually or spectroscopically by changes in absorption and/or emission properties.<ref name=":0">{{Cite book |last=Harris |first=Daniel C. |title=Exploring chemical analysis |date=2005 |publisher=W.H. Freeman |isbn=0-7167-0571-0 |edition=3rd |location=New York |oclc=54073810}}</ref> Hence, a pH indicator is a [[Chemical substance|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 [[color]] 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]]. The pH value of a neutral solution is 7.0 at 25°C ([[Standard conditions for temperature and pressure#Standard laboratory conditions|standard laboratory conditions]]). Solutions with a pH value below 7.0 are considered acidic and solutions with pH value above 7.0 are basic. Since most naturally occurring [[organic compound]]s are weak [[electrolyte]]s, such as [[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 (chemistry)|quantitative analysis]] of metal cations, the use of [[complexometric indicator]]s is preferred,<ref>{{cite book |last1=Schwarzenbach |first1=Gerold |translator-last=Irving |translator-first=Harry |title=Complexometric Titrations |date=1957 |publisher=[[Methuen & Co]] |location=London |pages=29–46 |edition=1st English}}</ref><ref>{{cite book |last1=West |first1=T.&nbsp;S. |title=Complexometry with EDTA and related reagents |date=1969 |publisher=[[BDH Chemicals|BDH Chemicals Ltd.]] |location=Poole, UK |pages=14–82 |edition=3rd}}</ref> whereas the third compound class, the [[redox indicator]]s, are used in [[redox titration]]s ([[titration]]s involving one or more redox reactions as the basis of chemical analysis).

==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.

==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 |url-access=subscription }}</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:#ffeb00" | yellow
| align="center" | 5.2
| align="center" | 6.8
| style="background:#5f00b9; color:white" | purple
|-
| [[Bromothymol blue]] 
| style="background:#ffd600" | yellow
| align="center" | 6.0
| align="center" | 7.6
| style="background:#114cff; 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
!Color
|-
|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 absorbances]], ''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.

== Equivalence point ==
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|date=2009|publisher=[[Houghton Mifflin Company]]|location=New York|pages=319–324|first=Steven S.|last=Zumdahl|edition=6th}}</ref>

== Naturally occurring pH indicators ==
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 and 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.

[[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.

''[[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.

Another natural pH indicator is the spice [[Turmeric#Indicator|turmeric]]. It turns yellow when exposed to [[acid]]s and reddish brown when in presence of an [[alkali]]s. 

{| class="wikitable"
! Indicator
! Low pH color
! High pH color
|-
|[[Hydrangea]] flowers
| style="background:blue; color:white" | blue
| style="background:fuchsia; color:white" | pink to purple
|-
|[[Anthocyanin]]s
| style="background:red; color:white" | red
| style="background:blue; color:white" | blue
|-
|[[Litmus]]
| style="background:red; color:white" | red
| style="background:blue; color:white" | blue
|-
|[[Turmeric]]
| style="background:yellow"|yellow
| style="background:brown; color:white" |reddish brown
|}
<gallery mode="packed" heights="120">
File:Blue Hydrangea.jpg|Hydrangea in acid soil
File:Hortensiapink.JPG|Hydrangea in alkaline soil
File:Indicateur chou rouge.jpg|A gradient of red cabbage extract pH indicator from acidic solution on the left to basic on the right
File:Purple Cauliflower Acid Base.jpg|Purple cauliflower soaked in baking soda (left) and vinegar (right). [[Anthocyanin]] acts as an pH indicator.
File:TurmericAcidBase.jpg|[[Turmeric]] dissolved in water is yellow under acidic and reddish brown under alkaline conditions
</gallery>

==See also==
* [[Chromophore]]
* [[Fecal pH test]]
* [[Nitrazine]]
* [[Sulfarsazene]]
* [[Universal indicator]]

==References==
{{reflist}}

==External links==
*{{Commonscatinline|pH indicators}}
*[http://www.ph-meter.info/pH-measurements-indicators Long indicator list], {{webarchive|url=https://web.archive.org/web/20220304003501/http://www.ph-meter.info/pH-measurements-indicators|date=4 March 2022}}

{{laboratory equipment}}
{{Authority control}}

[[Category:Chemical indicators]]
[[Category:Dyes]]
[[Category:Equilibrium chemistry]]
[[Category:Titration]]
[[Category:PH indicators|*]]