Indicator (chemistry)

pH indicator sticks

pH indicator paper for rough pH determination; left: lid with pH scale from 1 to 11; right: lower part of the jar with rolled up test strips

Indicators ( Latin indicare , "to show") are general aids that are intended to show certain information. They allow states to be determined and processes to be tracked by indicating that certain states have been reached or exited.

In chemistry , an indicator is understood to be a substance or a device that is used to monitor a chemical reaction or a state. Often a state is indicated by a certain color, the change by a color change. Indicators are most often used in titrations .

Depending on what kind of chemical or physical states or reactions one is tracking with an indicator, one differentiates between the following types of indicators:

pH indicators (for acid-base titrations and for determining the pH value)
Redox indicators (for redox titrations )
Complex indicators (in complexometry )
Thermal indicators (to display a temperature range)

pH indicators

An example that is often used for this is the indicator that shows the pH value of certain substances by comparing them with a color scale . Litmus , bromothymol blue or phenolphthalein are often used here . The latter only changes color when an alkaline solution is added . There are also universal indicators for a more precise measurement of the pH value . These are wetted with the solution to be tested, usually in the form of a paper strip, and the color of the paper strip is then compared with the color scale. One speaks now of either an acidic , alkaline or neutral solution .

The following table shows the colors of different indicators depending on the pH value .

pH indicators and their color scale

They are different

in the colors they have in acidic, neutral or alkaline solutions
in the pH range in which the change between the two colors takes place (this pH range is also called the transition range , which is pK _a ± 1).

Everyday acid-base indicators

Colors of red cabbage / blue cabbage juice, acidic on the left, alkaline on the right

Also cabbage juice can be used as a pH indicator. The dye cyanidin contained in red cabbage juice can take on colors from red = acidic to blue = alkaline (in an even more strongly alkaline environment it turns green and at pH> 10 it even turns yellow ). Depending on the method of preparation, red cabbage takes on a more red color (e.g. when adding apple pieces or a little vinegar ), or a purple to blue color (preparation with alkaline household baking soda ), which is what the common name red cabbage expresses.

Tea drinkers know how the tea color depends on the acid content: If lemon juice is added to the black tea , the color changes from dark brown to light reddish brown. This color change is also due to the coloring agents in the tea, which act as indicators.

functionality

The molecules of the acid-base indicators are themselves weak acids (or bases), that is, they can give up or take up protons . The indicator acid molecule is simply referred to as HInd. After releasing a proton, Ind ^- , the so-called corresponding indicator base, remains .

The protolysis equilibrium for the release or uptake of a proton in an aqueous indicator solution is the following:

{\ displaystyle \ mathrm {HInd _ {(aq)} + H_ {2} O _ {(l)} \ leftrightharpoons Ind _ {(aq)} ^ {-} + H_ {3} O _ {(aq)} ^ {+} }}

The indicator molecule can therefore release a proton but also take it up again. With a high concentration of H ₃ O ⁺ ions (i.e. in an acidic solution), the reaction to the left takes place to a greater extent, whereby the concentration of HInd (indicator acid) is greater than the concentration Ind ^- . With a very low concentration of H ₃ O ⁺ ions (i.e. in an alkaline solution), the reaction to the right takes place to a greater extent , which means that the concentration of Ind ^{- is} greater than the concentration of HInd.

This fact becomes particularly clear if one applies the law of mass action to the general equation mentioned above. Then with constant : ${\ displaystyle K_ {s}}$

{\ displaystyle K_ {s} = {c (\ mathrm {Ind} ^ {-}) \ cdot c (\ mathrm {H_ {3} O} ^ {+}) \ over c (\ mathrm {HInd})} }

As usual, the constant concentration of the water is included in the constant. It should be noted that the concentration of H ₃ O ^{+ is of} a different order of magnitude than that of the indicator. If this concentration is now greatly increased or decreased - depending on whether the solution becomes acidic or basic - the equilibrium has to be re-established, because K _s is constant. Therefore the concentration of Ind ^- has to change strongly, whereby at the same time, since HInd arises from Ind ^- , the concentration of HInd has to move correspondingly strongly in the other direction. Therefore, the indicator change is generally very quick.

The actual effect of the indicator is based on the fact that the connection HInd has a different color than Ind ^- . As a result of the protonation or deprotonation of the indicator, its mesomerism stabilization changes . In acidic solution, the concentration of HInd predominates, so that the solution takes on the color of the protonated form. If the pH value is increased, the concentration of Ind ^{- increases} , while the concentration of HInd decreases, until the former predominates and the solution takes on the color of Ind ^- (see e.g. phenolphthalein ). The different color of protonated and non-protonated dye molecules is known as halochromy .

The transition point of the indicator is characterized in that . At this point it also applies that the two terms in the above-mentioned equation of the law of mass action can be reduced, and . The pH value of the solution at the point of transition corresponds theoretically to the constant of the indicator. The changeover point perceptible to the human eye often deviates from this, since the human eye can only recognize the color change from a concentration ratio of 9: 1. The pH range that can be recognized as a mixed color is called the transition interval. ${\ displaystyle c (HInd) = c (Ind ^ {-})}$ ${\ displaystyle K_ {s} = c (H_ {3} O ^ {+})}$ ${\ displaystyle pK_ {s} = pH}$ ${\ displaystyle pK_ {s}}$

Indicator failure

During titration, the quantitatively precisely known addition of an acid or a base (titration agent) shifts the pH value of a buffer system to such an extent that the base or acid (Titrand) to be determined is completely neutralized. However, the pH indicator also represents a buffer system that simultaneously consumes hydroxide ions or oxonium ions from the titration agent. The concentration of indicators in the titration solution is usually in the order of 10 ⁻⁴ mol / l. The indicator hardly plays a role for much more concentrated titrands. In the analysis of natural water, however, the buffer concentration of the water constituents is the same to about 10 times higher, depending on the water hardness . Therefore, the indicator can cause a significant error here.

Another problem can arise because the indicator dye is usually added in the form of an alcoholic solution. This alone can change the properties of the entire buffer system.

Redox indicators

The simplest use of redox indicators is to determine the end point in redox titrations ( oxidimetry ).

Common redox indicators are:

Complex indicators (metal indicators)

A possible application is the dimensional analysis of the concentration of metal ions, for example complexometric titration. A typical application is the determination of water hardness .

Known complexometric indicators:

Solvatochrome indicators

Solvatochromic indicators change their color according to the surrounding solvent .

Voltage-dependent indicators

Voltage-dependent dyes change their color according to the electrical voltage

Thermal indicators (thermochromes)

Thermal indicators are often used where the temperature cannot simply be measured with a thermometer. For example, a melting pot is marked with thermal chalk and when a desired temperature has been reached in the flame, the color of the thermal chalk is displayed. Sufficient cooling can also be indicated by thermochromes.

Thermal indicators as stickers are based on liquid crystals . There are reversibly and irreversibly reacting variants. The latter are particularly suitable for registering the exceeding of certain temperature values at places invisible during operation.

Mixed indicator

Mixed indicators are mixtures of different indicators, whereby the turnover area is expanded or several turnover areas are generated. The mixed indicators also include the contrast indicators. A common example of a mixed indicator is Tashiro .

Contrast indicator

Contrast indicators usually consist of an indicator and a dye that retains its color. This addition increases the contrast of the envelope area. A common example of a contrast indicator is eriochrome black T mixed with methyl orange .

Humidity indicators

Blue gel dry (= blue)

Blue gel moist (= pink)

Humidity indicators change color when certain humidity levels to which they are exposed are exceeded. The color change reaction is based on water absorption, so the time of exposure also plays a role. The best known is the blue silica gel (blue gel) mixed with cobalt chloride , which turns violet or pink when exposed to moisture. It is applied to paper, added to moisture-sensitive consignments or housed in hermetically sealed assemblies behind inspection glasses. Due to the toxicity of the cobalt salt, cobalt-free alternatives are also offered ( orange gel ).

Fluorescent indicator

Fluorescence indicators are substances that change their fluorescence at the equivalence point of a titration . Also as fluorescent indicators are phosphors referred to, the stationary phase (separation layer) of plates for thin layer chromatography are admixed. They make it possible to recognize colorless substances under a UV lamp as a result of fluorescence quenching .

Individual evidence

^ E. Schweda: Jander / Blasius: Inorganic Chemistry I - Introduction & Qualitative Analysis . 17th edition. Hirzel, 2012, ISBN 978-3-7776-2134-0 , pp. 62 .
↑ Entry on thin layer chromatography. In: Römpp Online . Georg Thieme Verlag, accessed on June 18, 2014.

Web links

Prof. Blume's media offer for school chemistry: indicators

[1] E. Schweda: Jander / Blasius: Inorganic Chemistry I - Introduction & Qualitative Analysis . 17th edition. Hirzel, 2012, ISBN 978-3-7776-2134-0 , pp. 62 .

[2] Entry on thin layer chromatography. In: Römpp Online . Georg Thieme Verlag, accessed on June 18, 2014.