subject: Natural Acid Base Indicator [print this page] Acid base indicators are the substances which alter colors depending on the acidity or basicity of a testing solution. The acidity and basicity of a substance is measured in pH. While there are synthetic acid base indicators that are extensively used, there are also natural acid base indicators that can indicate a wide range of pH levels.
Examples of Natural Acid Base Indicator
Natural provides pH indicators in the form of plant pigments known as anthocyanins, these pigments change color over different pH ranges.
For example, there are indicators derived from red cabbage and red beet.
Red cabbage juice functions over a wide pH range from a low pH 1 up to pH 12.
Red beet juice will change from red color to yellow color between pH 11 and 12.
The indicator solution is prepared by extracting the juice of these substances.
There are other substances like tea and grapes which change colors when they react with acids or bases.
Many cyanidin-based compounds turn roses, apples, strawberries, autumn leaves, and cranberry juice into red. They make cornflowers, blueberries blue and some blackberries and some varieties of grapes purple.
Structural formula of Cyanidin:
Cynidin
Alizarin- Several synthetic modifications of alizarin in madder plant's root is also used as acid/base indicators.
Curcumin is a natural dye turmeric yellow found in curry powder. It turns colors from yellow at pH 7.4 to red at pH 8.6.
Structural formula of Anthocyanin
Anthocynin
It is a plant pigment that is a common acid/base indicator. Anthocyanin turns red cabbage purple, poppies red and cornflowers blue. Above are the change in color for the juice of red cabbage.
Conclusion for Natural Acid Base Indicators
Acid base indicators are the substances that change colors depending on the acidity or basicity of the solution. The colors are different for different pH levels. There are natural plant pigments available in shoots, flowers, fruits or roots or plants that can be extracted and used as acid base indicators.
A pH indicator (or pH paper) 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. Indicators can also show change in other physical properties; for example, olfactory indicators show change in their odor.
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 (alkaline). 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.
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 is in equilibrium with 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:
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.
