pH indicator
A pH indicator is a halochromic chemical compound that is added in small amounts to a solution so that the pH (acidity or basicity) of the solution can be determined visually. 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 colour of the solution to change depending on the pH. 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. As most naturally occurring organic compounds are weak protolytes, 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, whereas the third compound class, the redox indicators, are used in titrations involving a redox reaction as the basis of the analysis.
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Theory
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:
- HInd + H2O ⇌ H3O+ + Ind-
Here HInd stands for the acid form and Ind- 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:
![\textrm{pH} = \textrm{pK}_{a}%2B \log \frac{[\textrm{Ind}^-]}{[\textrm{HInd}]}](/2012-wikipedia_en_all_nopic_01_2012/I/9141ca3d4ac55eccf140dd2595acb375.png)
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.
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.
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 analytic chemistry and biology experiments to determine the extent of a chemical reaction. Because of the subjective determination of color, pH indicators are susceptible to imprecise readings. For applications requiring precise measurement of pH, a pH meter is frequently used.
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.
| Indicator | Low pH color | Transition pH range | High pH color |
|---|---|---|---|
| Gentian violet (Methyl violet 10B) | yellow | 0.0–2.0 | blue-violet |
| Leucomalachite green (first transition) | yellow | 0.0–2.0 | green |
| Leucomalachite green (second transition) | green | 11.6–14 | 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 |
| Bromophenol blue | yellow | 3.0–4.6 | purple |
| Congo red | blue-violet | 3.0–5.0 | red |
| Methyl orange | red | 3.1–4.4 | orange |
| Bromocresol green | yellow | 3.8–5.4 | blue |
| Methyl red | red | 4.4–6.2 | yellow |
| Methyl red | red | 4.5–5.2 | green |
| Azolitmin | 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 | colorless to reddish | 7.3–8.7 | greenish to blue |
| Cresol Red | yellow | 7.2–8.8 | reddish-purple |
| Phenolphthalein | colorless | 8.3–10.0 | fuchsia |
| Thymolphthalein | colorless | 9.3–10.5 | blue |
| Alizarine Yellow R | yellow | 10.2–12.0 | red |
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. [1]
Commercial preparations
Universal indicator and Hydrion papers are blends of different indicators that exhibits several smooth color changes over a very wide range of pH values.
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. Extracting anthocyanins from red cabbage leaves or the skin of a lemon to form a crude acid-base indicator is a popular introductory chemistry demonstration.
Anthocyanins can be extracted from a multitude of colored plants or plant parts, including from leaves (red cabbage); flowers (geranium, poppy, or rose petals); berries (blueberries, blackcurrant); and stems (rhubarb).
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' derives from 'coloured moss' in Old Norse. 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 aluminum available to these plants, turning the flowers blue. In alkaline soils, these reactions cannot occur and therefore aluminum is not taken up by the plant. As a result, the flowers remain pink.
| Indicator | Low pH color | High pH color |
|---|---|---|
| Hydrangea flowers | blue | pink to purple |
| Anthocyanins | red | blue |
| Litmus | red | blue |
See also
References
- ^ Steven S. Zumdahl (2009). Chemical Principles (6th ed.). New York: Houghton Mifflin Company. pp. 319-324.
- Long indicator list
- (French) Complete indicator listPDF (57.3 KiB)