Indicator Colors for pH Ranges

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This Demonstration displays the color changes of common pH indicators in some acidic and basic solutions. You can choose a strong acid, strong base, weak acid, or weak base. As you vary the molarity of the solution, the pH changes and the color of the chosen indicator changes accordingly. The value of the pH is also shown.

Contributed by: Isabella Iaccino and Gina Heinsohn (June 2016)
Special thanks to the University of Illinois NetMath program and the mathematics department at William Fremd High School.
Open content licensed under CC BY-NC-SA



Snapshot 1: a high-concentration solution of sodium hydroxide, which is a strong base, at a pH of 12.903 with bromcresol green as the indicator

Snapshot 2: a low-concentration solution of hydrochloric acid, which is a strong acid, at a pH of 1.495 with alizarin yellow as the indicator

Snapshot 3: a high-concentration solution of sodium bisulfate, which is a weak acid, at a pH of 1.597 with crystal violet as the indicator

For the pH to change accurately, there is a different formula for pH as molarity changes for each solution. The pH of hydrochloric acid, the strong acid, is equal to the negative log of the concentration of H+ ions. Because the reaction of hydrochloric acid with water is in a 1:1 ratio, the concentration of the H+ ions is equal to the molarity of the solution, so the pH of hydrochloric acid in terms of molarity is . The pH of sodium hydroxide, the strong base, is equal to , or . Finding a formula for a weak acid or weak base is more difficult, as the formula would be different for each weak acid or base based on the Ka/Kb value and the ratio of the balanced equation when it reacts with water. We set up an ICE problem (initial, change, equilibrium) and then used the equilibrium constant equations to solve for the final molarity of the hydronium/hydroxide ions in terms of the initial molarity of the substance. Then we substituted that value into the Henderson–Hasselbalch equation to find the pH/pOH. For the weak base we then subtracted the pOH value from 14 to get the pH value.

Each indicator changes colors within a different pH range. Outside of that range, they are either the initial or final color. The exception to this rule is phenolphthalein, which is colorless when outside of its range. The pH range and color changes for each indicator are as follows:

thymol blue, 1.2–2.8, red-yellow

alizarin yellow, 10.0–12.0, yellow-lilac

bromcresol green, 4.0–5.6, yellow-blue

phenol red, 6.4–8.0, yellow-red

phenolphthalein, 8.0–10.0, colorless-magenta-colorless

methyl green, 0.1–2.3, yellow-blue

crystal violet, 0.8–2.6, yellow-blue


[1] J. C. Kotz, P. M. Treichel, and J. R. Townsend, Chemistry & Chemical Reactivity, 8th ed. (Advanced Placement Edition), Belmont, CA: Brooks/Cole, Cengage Learning, 2012

[2] D. A. Skoog et al., Fundamentals of Analytical Chemistry, 8th ed., Belmont, CA: Thomson-Brooks/Cole, 2004.

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