A pH indicator or acid-base indicator is a chemical added in a small amount to a solution that causes a color change depending on the pH. This is a charge of common indicators, an explanation of how they work, and tips for choosing the right one for your needs.
How to Use a pH Indicator
An indicator doesn’t shift color at a precise pH or hydrogen ion concentration. Instead, the color change occurs over a range of hydrogen ion concentration. Titrate a weak acid using an indicator that changes under slightly alkaline conditions. Titrate a weak base using an indicator that changes color at a slightly acidic pH. When titrating strong acids or bases, aim for a pH indicator that displays a color change near a neutral pH.
Chart of Common pH Indicators
Here is a chart of common pH indicators, their pH range, their solutions, and their color changes. Some indicators display multiple color changes, so they occur on the list more than once. Various references list slightly different pH values and colors. This is because the pH range isn’t well-defined (expect accuracy within 1 pH value) and color is a judgement call.
|Indicator||pH Range||Quantity per 10 ml||Acid||Base|
|Thymol Blue||1.2-2.8||1-2 drops 0.1% soln. in aq.||red||yellow|
|Pentamethoxy red||1.2-2.3||1 drop 0.1% soln. in 70% alc.||red-violet||colorless|
|Tropeolin OO||1.3-3.2||1 drop 1% aq. soln.||red||yellow|
|2,4-Dinitrophenol||2.4-4.0||1-2 drops 0.1% soln. in 50% alc.||colorless||yellow|
|Methyl yellow||2.9-4.0||1 drop 0.1% soln. in 90% alc.||red||yellow|
|Methyl orange||3.1-4.4||1 drop 0.1% aq. soln.||red||orange|
|Bromophenol blue||3.0-4.6||1 drop 0.1% aq. soln.||yellow||blue-violet|
|Tetrabromophenol blue||3.0-4.6||1 drop 0.1% aq. soln.||yellow||blue|
|Alizarin sodium sulfonate||3.7-5.2||1 drop 0.1% aq. soln.||yellow||violet|
|α-Naphthyl red||3.7-5.0||1 drop 0.1% soln. in 70% alc.||red||yellow|
|p-Ethoxychrysoidine||3.5-5.5||1 drop 0.1% aq. soln.||red||yellow|
|Bromocresol green||4.0-5.6||1 drop 0.1% aq. soln.||yellow||blue|
|Methyl red||4.4-6.2||1 drop 0.1% aq. soln.||red||yellow|
|Bromocresol purple||5.2-6.8||1 drop 0.1% aq. soln.||yellow||purple|
|Chlorophenol red||5.4-6.8||1 drop 0.1% aq. soln.||yellow||red|
|Bromophenol blue||6.2-7.6||1 drop 0.1% aq. soln.||yellow||blue|
|p-Nitrophenol||5.0-7.0||1-5 drops 0.1% aq. soln.||colorless||yellow|
|Azolitmin||5.0-8.0||5 drops 0.5% aq. soln.||red||blue|
|Phenol red||6.4-8.0||1 drop 0.1% aq. soln.||yellow||red|
|Neutral red||6.8-8.0||1 drop 0.1% soln. in 70% alc.||red||yellow|
|Rosolic acid||6.8-8.0||1 drop 0.1% soln. in 90% alc.||yellow||red|
|Cresol red||7.2-8.8||1 drop 0.1% aq. soln.||yellow||red|
|α-Naphtholphthalein||7.3-8.7||1-5 drops 0.1% soln. in 70% alc.||rose||green|
|Tropeolin OOO||7.6-8.9||1 drop 0.1% aq. soln.||yellow||rose-red|
|Thymol blue||8.0-9.6||1-5 drops 0.1% aq. soln.||yellow||blue|
|Phenolphthalein||8.0-10.0||1-5 drops 0.1% soln. in 70% alc.||colorless||red|
|α-Naphtholbenzein||9.0-11.0||1-5 drops 0.1% soln. in 90% alc.||yellow||blue|
|Thymolphthalein||9.4-10.6||1 drop 0.1% soln. in 90% alc.||colorless||blue|
|Nile blue||10.1-11.1||1 drop 0.1% aq. soln.||blue||red|
|Alizarin yellow||10.0-12.0||1 drop 0.1% aq. soln.||yellow||lilac|
|Salicyl yellow||10.0-12.0||1-5 drops 0.1% soln. in 90% alc.||yellow||orange-brown|
|Diazo violet||10.1-12.0||1 drop 0.1% aq. soln.||yellow||violet|
|Tropeolin O||11.0-13.0||1 drop 0.1% aq. soln.||yellow||orange-brown|
|Nitramine||11.0-13.0||1-2 drops 0.1% soln in 70% alc.||colorless||orange-brown|
|Poirrier’s blue||11.0-13.0||1 drop 0.1% aq. soln.||blue||violet-pink|
|Trinitrobenzoic acid||12.0-13.4||1 drop 0.1% aq. soln.||colorless||orange-red|
In addition to the pH indicators on this list, there are many natural acid-base indicators you can make using fruits, vegetables, flowers, juices, and spices. Red or purple cabbage juice is the best-known of these.
Universal indicator is a mixture of several different pH indicators that displays smooth color changes over a range of pH values. There are multiple universal indicator formulas, so the pH ranges and colors depend on the formula. The most common ones are variations of Yamada’s formula, published in 1933. A typical recipe includes 1-propanol, sodium salt, sodium hydroxide, monosodium salt, phenolphthalein, methyl red, bromothymol blue, and thymol blue. This mixture displays the colors red, orange-yellow, green, blue, and indigo-violet:
|< 3||Red||Strongly acidic|
|3 to 6||Orange to Yellow||Weakly acidic|
|8 to 11||Blue||Weakly alkaline (basic)|
|> 11||Indigo to Violet||Strongly alkaline (basic)|
How pH Indicators Work
Most pH indicators are weak acids or weak bases. They dissociate according to the general chemical reaction:
HInd + H2O ⇌ H3O+ + Ind−
Here, HInd is the acid form of the indicator and Ind− is its conjugate base. The ratio between HInd and Ind− determines the color of the solution and indirectly indicates the pH of the solution according to the Henderson-Hasselbalch equation:
pH = pKa + log10 [Ind−] / [HInd]
Keep in mind, the color change of a pH indicator isn’t instantaneous. Instead, there is a pH range where a mix of the acid and conjugate base colors appear. An indicator gives a reasonably accurate pH value within a pH or pKa value of plus or minus one.
How to Choose a pH Indicator
The most important step in choosing the right pH indicator is picking one that has a color change within the pH range of the chemical reaction being studied. For a titration, you ideally want to pick a pH indicator that changes color right at the equivalence point. In practice, it’s almost impossible to find an indicator that changes color at the exact pH value, so you have to go with one that changes color at a slightly higher or lower pH. In this case, you titrate until you see the color change closest to the equivalence point.
For example, say you titrate a strong acid to a strong base. The equivalence point for this reaction is a pH value of 7. If you use phenolphthalein, you expect a pink/red color to vanish below a pH value of 8.0. You titrate until the solution becomes colorless because that is as close to the equivalence point as you can get. If you use methyl orange, you expect the color to change from yellow to orange somewhere below a pH of 6 and from orange to red around a pH of 4. For the strong acid to strong base reaction, titrate until yellow just starts to turn orange. If you wait until the color changes to red, you’re way past the equivalence point.
If you can choose between indicators that change color at the desired pH, go with the one that shows the sharpest color change. For example, bromophenol blue and p-nitrophenol both give color changes around a neutral pH, but the change from yellow to blue (bromophenol) is easier to see than the one from colorless to yellow (p-nitrophenol).
Other factors than matter include solvent (alcohol or water), price, and versatility. Which pH indicator you choose is a matter of its pH range, color change, solvent, availability, and cost.
- Foster, L. S.; Gruntfest, I. J. (1937). “Demonstration experiments using universal indicators”. Journal of Chemical Education. 14 (6): 274. doi:10.1021/ed014p274
- Lange, Norbert A. (1952). Lange’s Handbook of Chemistry (8th ed.). Handbook Publishers Inc. ASIN : B000RFWWKO
- Kolthoff, I. M.; Stenger, V. A. (1942). Volumetric Analysis. Interscience Publishers, Inc., New York. ISBN: 978-0470500507
- Schwarzenbach, Gerold (1957). Complexometric Titrations. Translated by Irving, Harry. London: Methuen & Co.
- Zumdahl, Steven S. (2009). Chemical Principles (6th ed.). New York: Houghton Mifflin Company. ISBN: 978-0618946907.