PH indicator

Last updated December 05, 2024

A pH indicator is a halochromic chemical compound added in small amounts to a solution so the pH (acidity or basicity) of the solution can be determined visually or spectroscopically by changes in absorption and/or emission properties.^[1] Hence, a pH indicator is a chemical detector for hydronium ions (H₃O⁺) or hydrogen ions (H⁺) in the 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 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 compounds are weak electrolytes, such as carboxylic acids and amines, 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 indicators is preferred,^[2]^[3] whereas the third compound class, the redox indicators, are used in redox titrations (titrations 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_(aq) + H^₂O_(l)⇌H^₃O⁺_(aq) + Ind⁻_(aq)

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

Vice versa for basic pH indicators in aqueous solutions:

IndOH_(aq) + H^₂O_(l)⇌H^₂O_(l) + Ind⁺_(aq) + OH⁻_(aq)

where "IndOH" stands for the basic form and "Ind⁺" 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 = pK_a + log₁₀ ⁠ [Ind⁻] [HInd] ⁠

or

pOH = pK_b + log₁₀⁠ [Ind⁺] [IndOH] ⁠

The equations, derived from the acidity constant and basicity constant, states that when pH equals the pK_a or pK_b value of the indicator, both species are present in a 1:1 ratio. If pH is above the pK_a or pK_b 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 pK_a or pK_b value, the converse is true.

Usually, the color change is not instantaneous at the pK_a or pK_b 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 pK_a or pK_b 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 pK_a + 1 or pK_b + 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 pK_a − 1 or pK_b − 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

pH indicators are frequently employed in titrations in analytical chemistry and biology to determine the extent of a chemical reaction.^[1] Because of the 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 papers) 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.^[1]

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.

Indicator	Low pH color	Transition low end	Transition high end	High pH color
Gentian violet (Methyl violet 10B)^[4]	yellow	0.0	2.0	blue-violet
Malachite green (first transition)	yellow	0.0	2.0	green
Malachite green (second transition)	green	11.6	14.0	colorless
Thymol blue (first transition)	red	1.2	2.8	yellow
Thymol blue (second transition)	yellow	8.0	9.6	blue
Methyl yellow	red	2.9	4.0	yellow
Methylene blue	colorless	5.0	9.0	dark blue
Bromophenol blue	yellow	3.0	4.6	blue
Congo red	blue-violet	3.0	5.0	red
Methyl orange	red	3.1	4.4	yellow
Screened methyl orange (first transition)	red	0.0	3.2	purple-grey
Screened methyl orange (second transition)	purple-grey	3.2	4.2	green
Bromocresol green	yellow	3.8	5.4	blue
Methyl red	red	4.4	6.2	yellow
Methyl purple	purple	4.8	5.4	green
Azolitmin (litmus)	red	4.5	8.3	blue
Bromocresol purple	yellow	5.2	6.8	purple
Bromothymol blue	yellow	6.0	7.6	blue
Phenol red	yellow	6.4	8.0	red
Neutral red	red	6.8	8.0	yellow
Naphtholphthalein	pale red	7.3	8.7	greenish-blue
Cresol red	yellow	7.2	8.8	reddish-purple
Cresolphthalein	colorless	8.2	9.8	purple
Phenolphthalein (first transition)	colorless	8.3	10.0	purple-pink
Phenolphthalein (second transition)	purple-pink	12.0	13.0	colorless
Thymolphthalein	colorless	9.3	10.5	blue
Alizarine Yellow R	yellow	10.2	12.0	red
Indigo carmine	blue	11.4	13.0	yellow

Universal Indicator

pH range	Description	Colour
1-3	Strong acid	Red
3 – 6	Weak acid	Orange/Yellow
7	Neutral	Green
8 – 11	Weak alkali	Blue
11-14	Strong alkali	Violet/Indigo

Precise pH measurement

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⁺ and A⁻.

HA ⇌ H⁺ + A⁻

The value of the acid dissociation constant, pK_a, must be known. The molar absorbances, ε_HA and ε_A⁻ of the two species HA and A⁻ at wavelengths λ_x and λ_y must also have been determined by previous experiment. Assuming Beer's law to be obeyed, the measured absorbances A_x and A_y at the two wavelengths are simply the sum of the absorbances due to each species.

{\begin{aligned}A_{x}&=[{\ce {HA}}]\varepsilon _{{\ce {HA}}}^{x}+[{\ce {A-}}]\varepsilon _{{\ce {A-}}}^{x}\\A_{y}&=[{\ce {HA}}]\varepsilon _{{\ce {HA}}}^{y}+[{\ce {A-}}]\varepsilon _{{\ce {A-}}}^{y}\end{aligned}}

These are two equations in the two concentrations [HA] and [A⁻]. Once solved, the pH is obtained as

\mathrm {pH} =\mathrm {p} K_{\mathrm {a} }+\log {\frac {[{\ce {A-}}]}{[{\ce {HA}}]}}

If measurements are made at more than two wavelengths, the concentrations [HA] and [A⁻] can be calculated by 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.

When a single indicator is used, this method is limited to measurements in the pH range pK_a ± 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.^[5]

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 (geranium, poppy, or rose petals); berries (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. It turns yellow when exposed to acids and reddish brown when in presence of an alkalis.

Indicator	Low pH color	High pH color
Hydrangea flowers	blue	pink to purple
Anthocyanins	red	blue
Litmus	red	blue
Turmeric	yellow	reddish brown

Hydrangea in acid soil
Hydrangea in alkaline soil
A gradient of red cabbage extract pH indicator from acidic solution on the left to basic on the right
Purple cauliflower soaked in baking soda (left) and vinegar (right). Anthocyanin acts as an pH indicator.
Turmeric dissolved in water is yellow under acidic and reddish brown under alkaline conditions

Related Research Articles

An acid is a molecule or ion capable of either donating a proton (i.e. hydrogen ion, H⁺), known as a Brønsted–Lowry acid, or forming a covalent bond with an electron pair, known as a Lewis acid.

Analytical chemistry studies and uses instruments and methods to separate, identify, and quantify matter. In practice, separation, identification or quantification may constitute the entire analysis or be combined with another method. Separation isolates analytes. Qualitative analysis identifies analytes, while quantitative analysis determines the numerical amount or concentration.

In chemistry, an acid–base reaction is a chemical reaction that occurs between an acid and a base. It can be used to determine pH via titration. Several theoretical frameworks provide alternative conceptions of the reaction mechanisms and their application in solving related problems; these are called the acid–base theories, for example, Brønsted–Lowry acid–base theory.

In chemistry, pH, also referred to as acidity or basicity, historically denotes "potential of hydrogen". It is a logarithmic scale used to specify the acidity or basicity of aqueous solutions. Acidic solutions are measured to have lower pH values than basic or alkaline solutions.

Titration is a common laboratory method of quantitative chemical analysis to determine the concentration of an identified analyte. A reagent, termed the titrant or titrator, is prepared as a standard solution of known concentration and volume. The titrant reacts with a solution of analyte to determine the analyte's concentration. The volume of titrant that reacted with the analyte is termed the titration volume.

A conjugate acid, within the Brønsted–Lowry acid–base theory, is a chemical compound formed when an acid gives a proton to a base—in other words, it is a base with a hydrogen ion added to it, as it loses a hydrogen ion in the reverse reaction. On the other hand, a conjugate base is what remains after an acid has donated a proton during a chemical reaction. Hence, a conjugate base is a substance formed by the removal of a proton from an acid, as it can gain a hydrogen ion in the reverse reaction. Because some acids can give multiple protons, the conjugate base of an acid may itself be acidic.

A buffer solution is a solution where the pH does not change significantly on dilution or if an acid or base is added at constant temperature. Its pH changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. In nature, there are many living systems that use buffering for pH regulation. For example, the bicarbonate buffering system is used to regulate the pH of blood, and bicarbonate also acts as a buffer in the ocean.

In chemistry, an acid dissociation constant is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction

Solubility equilibrium is a type of dynamic equilibrium that exists when a chemical compound in the solid state is in chemical equilibrium with a solution of that compound. The solid may dissolve unchanged, with dissociation, or with chemical reaction with another constituent of the solution, such as acid or alkali. Each solubility equilibrium is characterized by a temperature-dependent solubility product which functions like an equilibrium constant. Solubility equilibria are important in pharmaceutical, environmental and many other scenarios.

In chemistry, there are three definitions in common use of the word "base": Arrhenius bases, Brønsted bases, and Lewis bases. All definitions agree that bases are substances that react with acids, as originally proposed by G.-F. Rouelle in the mid-18th century.

<span class="mw-page-title-main">Phenolphthalein</span> pH indicator that turns pink in basic solution

Phenolphthalein ( feh-NOL(F)-thə-leen) is a chemical compound with the formula C₂₀H₁₄O₄ and is often written as "HIn", "HPh", "phph" or simply "Ph" in shorthand notation. Phenolphthalein is often used as an indicator in acid–base titrations. For this application, it turns colorless in acidic solutions and pink in basic solutions. It belongs to the class of dyes known as phthalein dyes.

In chemistry and biochemistry, the Henderson–Hasselbalch equation $relates the pH of a chemical solution of a weak acid to the numerical value of the acid dissociation constant, K a, of acid and the ratio of the concentrations, of the acid and its conjugate base in an equilibrium.$

<span class="mw-page-title-main">Neutralization (chemistry)</span> Chemical reaction in which an acid and a base react quantitatively

In chemistry, neutralization or neutralisation is a chemical reaction in which acid and a base react with an equivalent quantity of each other. In a reaction in water, neutralization results in there being no excess of hydrogen or hydroxide ions present in the solution. The pH of the neutralized solution depends on the acid strength of the reactants.

An acid–base titration is a method of quantitative analysis for determining the concentration of Brønsted-Lowry acid or base (titrate) by neutralizing it using a solution of known concentration (titrant). A pH indicator is used to monitor the progress of the acid–base reaction and a titration curve can be constructed.

The Brønsted–Lowry theory (also called proton theory of acids and bases) is an acid–base reaction theory which was first developed by Johannes Nicolaus Brønsted and Thomas Martin Lowry independently in 1923. The basic concept of this theory is that when an acid and a base react with each other, the acid forms its conjugate base, and the base forms its conjugate acid by exchange of a proton (the hydrogen cation, or H⁺). This theory generalises the Arrhenius theory.

<span class="mw-page-title-main">Methyl orange</span> Chemical compound

Methyl orange is a pH indicator frequently used in titration because of its clear and distinct color variance at different pH values. Methyl orange shows red color in acidic medium and yellow color in basic medium. Because it changes color at the pK_a of a mid strength acid, it is usually used in titration of strong acids in weak bases that reach the equivalence point at a pH of 3.1-4.4. Unlike a universal indicator, methyl orange does not have a full spectrum of color change, but it has a sharp end point. In a solution becoming less acidic, methyl orange changes from red to orange and, finally, to yellow—with the reverse process occurring in a solution of increasing acidity.

Acid salts are a class of salts that produce an acidic solution after being dissolved in a solvent. Its formation as a substance has a greater electrical conductivity than that of the pure solvent. An acidic solution formed by acid salt is made during partial neutralization of diprotic or polyprotic acids. A half-neutralization occurs due to the remaining of replaceable hydrogen atoms from the partial dissociation of weak acids that have not been reacted with hydroxide ions to create water molecules.

Equilibrium constants are determined in order to quantify chemical equilibria. When an equilibrium constant $K$ is expressed as a concentration quotient,

Conductometry is a measurement of electrolytic conductivity to monitor a progress of chemical reaction. Conductometry has notable application in analytical chemistry, where conductometric titration is a standard technique. In usual analytical chemistry practice, the term conductometry is used as a synonym of conductometric titration while the term conductimetry is used to describe non-titrative applications. Conductometry is often applied to determine the total conductance of a solution or to analyze the end point of titrations that include ions.

Acid strength is the tendency of an acid, symbolised by the chemical formula $, to dissociate into a proton,, and an anion, . The dissociation or ionization of a strong acid in solution is effectively complete, except in its most concentrated solutions.$

References

1 2 3 Harris, Daniel C. (2005). Exploring chemical analysis (3rd ed.). New York: W.H. Freeman. ISBN 0-7167-0571-0. OCLC 54073810.
↑ Schwarzenbach, Gerold (1957). Complexometric Titrations. Translated by Irving, Harry (1st English ed.). London: Methuen & Co. pp. 29–46.
↑ West, T. S. (1969). Complexometry with EDTA and related reagents (3rd ed.). Poole, UK: BDH Chemicals Ltd. pp. 14–82.
↑ Adams, Elliot Q.; Rosenstein, Ludwig. (1914). "The Color and Ionization of Crystal-Violet". Journal of the American Chemical Society. 36 (7): 1452–1473. doi:10.1021/ja02184a014. hdl: 2027/uc1.b3762873 . ISSN 0002-7863.
↑ Zumdahl, Steven S. (2009). Chemical Principles (6th ed.). New York: Houghton Mifflin Company. pp. 319–324.

External links

Media related to pH indicators at Wikimedia Commons
Long indicator list, Archived 4 March 2022 at the Wayback Machine

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[:0-1] 1 2 3 Harris, Daniel C. (2005). Exploring chemical analysis (3rd ed.). New York: W.H. Freeman. ISBN 0-7167-0571-0. OCLC 54073810.

[2] Schwarzenbach, Gerold (1957). Complexometric Titrations. Translated by Irving, Harry (1st English ed.). London: Methuen & Co. pp. 29–46.

[3] West, T. S. (1969). Complexometry with EDTA and related reagents (3rd ed.). Poole, UK: BDH Chemicals Ltd. pp. 14–82.

[4] Adams, Elliot Q.; Rosenstein, Ludwig. (1914). "The Color and Ionization of Crystal-Violet". Journal of the American Chemical Society. 36 (7): 1452–1473. doi:10.1021/ja02184a014. hdl: 2027/uc1.b3762873 . ISSN 0002-7863.

[5] Zumdahl, Steven S. (2009). Chemical Principles (6th ed.). New York: Houghton Mifflin Company. pp. 319–324.

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