An indicator is used to indicate the end of the titration as a noticeable pH change occurs near the equivalence point of the acid-base titration. While choosing an indicator for acid-base titration, select an indicator whose pH range falls within the pH change of the reaction.
Compounds that change color when exposed to acidic or alkaline solutions are known as indicators. Colored indicators are widely used to detect pH and can be added to the reaction mixture to determine the titration's endpoint or equivalence point.
In the laboratory of science classes for practical purposes, indicators like phenolphthalein and methyl orange are commonly used. A universal indicator, pH paper, litmus paper (Blue or Red), and pH meter are usually used to determine the pH of the substance.
In the titration process, a substance to be analyzed (Titrand) is placed in a conical flask, two to four drops of a suitable indicator are added, followed by a drop by drop addition of titrant of known strength from a burette until the completion of a chemical reaction.
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History of indicators:
Sir Robert Boyle first described the use of a natural dye as an acid-base indicator in 1664 in his collection of articles "The Experimental History of Dyes." By using markers for the experimental classification of acids and bases, Boyle made significant contributions to the early theory of acids and bases.
The theory of acid-base indicators can be explained in two ways: Ostwald's theory and the Quinonoid theory. In chemistry, there are three types of indicators: natural indicators, artificial indicators, and olfactory indicators, however artificial and natural indicators are the two most commonly used.
What are the two criteria for an indicator to be used in titration?
The two general criteria for an indicator to be used in a titration are: the pH at the end of the titration should be close to the indicator's neutral point, and the indicator should indicate a sharp color change near the equivalence point of the titration point.
For example, when titrating a strong acid with a strong base, the pH quickly shifts from 03 to 11. Therefore, the phenolphthalein indicator, which has a pH range of 08 to 10, is a good choice for this type of titration.
Choice of indicators for various acid-base titrations:
Each pH indicator changes color over a defined range of pH, known as the indicator range. The indicator ranges are listed below for some of the indicators.
Titration of a strong acid against a strong base:
For example, hydrochloric acid (HCl) vs. sodium hydroxide (NaOH)
In this form of titration, the pH value at the endpoint changes about from 04 to 10. As a result, any indicator that changes color within this range can be used to titrate strong acid against a strong base. For this type of titration, phenolphthalein may be suitable as an indicator.
Titration of a weak acid against a strong base:
For example, Oxalic acid vs. sodium hydroxide (NaOH)
In this type of titration, the pH value changes slightly at the endpoint which is about 6.5 to 10. Since the working range of phenolphthalein is 8.3 to 10, it is a suitable indicator for this titration. In these cases, methyl orange is not an appropriate indicator, because it has a working range below pH 05.
Titration of strong acid against weak base:
For example, hydrochloric acid (HCl) vs. sodium carbonate (Na2CO3)
When HCl (strong acid) is titrated against sodium carbonate (a weak base), at the endpoint the pH varies from 3.5 to 7.5. Methyl orange, which changes color within this pH range, is the more suitable indicator for this type of titration.
Titration of weak acid against weak base:
For example, Acetic acid (CH3COOH) vs. ammonium hydroxide or ammonia solution (NH4OH)
pH value at the end of this form of titration there is no sharp change. Therefore, this no indicator is suitable, because it requires a vertical portion of the curve over two pH units. As a result, neither phenolphthalein nor methyl orange is useful.
Indicators for different types of titration:
There are four types of titration such as acid-base titrations, redox titrations, precipitation titrations, and complexometric titrations, etc. According to the pH range and reaction, each titration has a particular reaction mechanism and uses different indicators.
Indicators for acid-base titration:
Acid-base indicators are generally classified into listed three groups: phthaleins and sulphophthaleins, Azo indicators, Triphenylmethane indicators.
Examples: Phenolphthalein indicator, methyl orange indicator, and malachite green indicator, etc. can be used be as an indicator
Indicators for complexometric titration:
Complexometric indicators are water-soluble organic compounds that are used to detect the endpoint of a titration.
Examples: Eriochrome Black T, calgamite, arsenazo, xylenol orange, Eriochrome red B, fast sulphon black, and calcein, etc. can be used be as an indicator
Indicators for redox titration:
The redox indicators are indicators that show a reversible color change between oxidized and reduced forms that change color at a specific potential.
Examples: Potassium permanganate (self-indicator), phenanthroline blue, methylene blue, 1, 10 phenanthroline monohydrate, and safrannin-T, etc. can be used be as an indicator
Indicators for precipitation titration:
There are three types of indicators employed in precipitation titration: two of them are chloride ions, and one is cation silver.
Examples: In Mohr's titration, silver nitrate or potassium chromate can be used, In Volhard’s titration, the ferric ion can use, and In Fajan’s titration, the dye dichlorofluorescein can used be as an indicator.
Indicators for non-aqueous titration:
The ionized and unionized forms of indicators, as well as numerous resonant forms of indicators, are all generally useful for non-aqueous titration.
Examples: Crystal violet indicator, methyl red indicator, naphtholbenzein indicator, Quinaldine red, and thymol blue, etc. can be used be as an indicator
List of indicators:
The following table gives an approximate color of some indicators at different pH values, as well as the types of titrations for which they are useful.
| Name of Indicator | Range of pH | Colour in acidic | Colour in basic |
|---|---|---|---|
| Methyl Violet | 0.0 to 1.6 | Yellow | Blue |
| Crystal Violet | 0.0 to 1.8 | Yellow | Blue |
| Cresol red | 0.2 to 1.8 | Yellow | Red |
| Malachite green | 0.2 to 1.8 | Yellow | Green |
| Thymol Blue | 1.2 to 2.8 | Red | Yellow |
| Methyl Yellow | 2.9 to 4.0 | Red | Yellow |
| Bromophenol blue | 3.0 to 4.6 | Yellow | Blue |
| Methyl Orange | 3.1 to 4.4 | Red | Yellow |
| Bromocresol Green | 3.8 to 5.4 | Yellow | Blue |
| Dichlorofluorescein | 4.0 to 6.6 | Colorless | Green |
| Methyl Red | 4.2 to 6.2 | Red | Yellow |
| Chlorophenol Red | 4.8 to 6.4 | Yellow | Red |
| Bromocresol Purple | 5.2 to 6.6 | Yellow | Purple |
| Bromothymol Blue | 6.0 to 7.6 | Yellow | Blue |
| Phenol Red | 6.4 to 8.0 | Yellow | Red |
| Cresol Purple | 7.4 to 9.0 | Yellow | Purple |
| Thymol Blue | 8.0 to 9.6 | Yellow | Blue |
| Phenolphthalein | 8.0 to 9.8 | Colorless | Red |
| Thymolphthalein | 9.3 to 10.5 | Colorless | Blue |
| Alizarin Yellow | 10.1 to 12.0 | Yellow | Red |
| Indigo Carmine | 11.4 to 13.0 | Blue | Yellow |
| - | - | - | - |
| Litmus | 5.0 to 8.0 | Blue | Red |
| Universal indicator | 0.1 to 14.0 | - | - |
FAQ (Frequently Asked Questions):
Why is universal indicator not used in titration?
A universal indicator is used to determine the approximate pH of a sample compound. As it changes color over a broad range of pH it is not used in the titration.
What is a natural and artificial indicator? Give some examples.
Natural indicators are types of indicators that occur naturally (e.g turmeric, red cabbage) and artificial indicators are those that are prepared artificially in the lab or produced through a chemical reaction (e.g. phenolphthalein, methyl orange), both can be used to identify whether a substance is acidic or alkaline.
What happens if you use the wrong indicator in titration?
If we use the wrong indicator in a titration, the endpoint will not be as expected, as the titrant solution may take more or less volume, making your calculations incorrect. As a result, the result may be wrong or incorrect.
References:
- Reactions of Acids and Bases in Analytical Chemistry. Hulanicki, A. and Masson, M.R. New York: Halsted Press, 1987.
- https://www.sciencecompany.com/ph_indicator_ranges.aspx
- https://en.wikipedia.org/wiki/Titration
- Aqueous Acid-Base Equilibria and Titrations. Levie, Robert De. New York : Oxford University Press, 1999.
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