An acid-base indicator is a substance that changes color depending on whether a solution is acidic or basic. These indicators are typically weak organic acids or bases themselves, and their color shifts happen because the molecule physically changes form as the pH of the surrounding solution rises or falls. You’ve probably seen one in action if you’ve ever used litmus paper or watched a chemistry demonstration where a clear liquid suddenly turned pink.
How Indicators Work
The color change comes down to a simple chemical tug-of-war. An indicator molecule exists in two forms: one that has an extra hydrogen ion attached (the “protonated” form) and one that has lost that hydrogen ion (the “deprotonated” form). Each form absorbs light differently, so each form looks like a different color to your eye.
When you drop an indicator into an acidic solution, the abundance of hydrogen ions in that solution pushes the indicator toward its protonated form, producing one color. Add a base, and the hydrogen ions get pulled away, shifting the indicator toward its deprotonated form and a second color. The solution isn’t doing anything magical. It’s just pushing the indicator molecule back and forth between two versions of itself, and each version happens to look different.
Take methyl orange as an example. In its protonated form, it’s red. Lose that hydrogen ion, and it turns yellow. So in a strongly acidic solution, methyl orange looks red. As you add base and neutralize the acid, it gradually shifts to yellow.
The Transition Range
Indicators don’t flip from one color to another at a single exact pH. Instead, the change happens gradually across a range of about 2 pH units. This window is called the transition range, and it’s centered roughly around a value chemists call the indicator’s pKa, which is the pH at which exactly half the indicator molecules are in one form and half are in the other.
Below the transition range, virtually all the indicator is in one color form. Above it, virtually all is in the other. In between, you see a blend. For phenolphthalein, the transition range runs from about pH 8.3 to 10.5. Below 8.3 the solution looks colorless; above 10.5 it looks pink to red. For methyl orange, the range is pH 3.1 to 4.4, shifting from red to orange-yellow. Bromothymol blue changes from yellow to blue between pH 5.8 and 7.6, passing through green at neutral pH.
This is why no single indicator works for every situation. Each one is tuned to a specific slice of the pH scale.
Common Indicators and Their Colors
- Phenolphthalein: Colorless in acidic and neutral solutions, turning pink to red above about pH 8.3. One of the most recognizable indicators in chemistry labs.
- Methyl orange: Red below pH 3.1, transitioning to orange-yellow by pH 4.4. Useful for detecting strongly acidic conditions.
- Bromothymol blue: Yellow in acidic solutions, green around neutral pH, and blue in basic solutions. Its transition range (pH 5.8 to 7.6) makes it handy for detecting changes near neutral.
- Litmus: Red in acids, blue in bases. Litmus is one of the oldest indicators, derived from lichens, and works broadly rather than at a precise pH.
Indicators in Titrations
The most common practical use for acid-base indicators is in titrations, where you slowly add a base to an acid (or vice versa) to figure out the concentration of one of them. The indicator tells you when to stop adding.
There’s an important distinction here. The equivalence point is the exact moment when the acid and base have completely neutralized each other, with no excess of either remaining. The endpoint is the moment the indicator actually changes color in front of you. These two points are close, but not always identical. The endpoint typically falls just after the equivalence point, making it a very good approximation rather than a perfect match.
Choosing the right indicator for a titration matters. You want an indicator whose transition range overlaps with the pH at the equivalence point. For a strong acid mixed with a strong base, the equivalence point lands right around pH 7, so bromothymol blue (which shifts color near pH 7) works well. For a weak acid titrated with a strong base, the equivalence point sits higher on the pH scale, often around pH 8 to 10, making phenolphthalein the better choice. Using methyl orange for that same weak-acid titration would cause it to change color too early, well before neutralization was actually complete, and your results would be off.
Natural Indicators
You don’t need a chemistry supply catalog to find an acid-base indicator. Many plants produce pigments that respond to pH changes, most notably a group of compounds called anthocyanins. Red cabbage is the classic example. Chop it up, boil it in water, and the purple liquid you get shifts to red or pink in acids and turns green or yellow in bases. It covers a remarkably wide pH range, which is why it shows up in so many science fair projects.
Turmeric, the bright yellow spice made from the root of the Curcuma longa plant, also works as an indicator. It stays yellow in acidic and neutral solutions but turns reddish-brown in basic ones. Beetroot juice, grape juice, and extracts from China rose petals all behave similarly, each with their own color shifts. These natural indicators work on the same principle as synthetic ones: the pigment molecule changes structure in response to gaining or losing hydrogen ions, and that structural change alters which wavelengths of light it absorbs.
Why the Color Change Happens
At a deeper level, what’s actually changing is the shape of the indicator molecule’s electron cloud. When a hydrogen ion attaches to or detaches from the molecule, it rearranges the way electrons are distributed across the molecule’s structure. Since the color of any substance depends on which wavelengths of light its electrons absorb, shifting those electrons around changes the color you see. A small chemical event, gaining or losing a single hydrogen ion, produces a visible, dramatic result. That’s what makes indicators so useful: they turn an invisible chemical process into something you can see with the naked eye.