Indicators change color in titration because they are weak acids or bases themselves, and each form (acidic or basic) has a different color. When the solution’s pH shifts past a certain point during titration, the indicator molecules switch from one form to the other, producing a visible color change that signals the endpoint.
Understanding exactly how this works helps you pick the right indicator, read your results accurately, and avoid common mistakes in the lab.
How Indicator Molecules Work
An acid-base indicator is a molecule that exists in two forms: a protonated form (with an extra hydrogen ion attached) and a deprotonated form (with that hydrogen ion removed). These two forms absorb light at different wavelengths, so they appear as different colors to your eye. Phenolphthalein, for example, is colorless when protonated and pink when deprotonated.
In acidic solution, there are plenty of hydrogen ions floating around. They bind to the indicator, keeping it in its protonated form and displaying one color. As you add base during a titration, you neutralize those hydrogen ions. Eventually the solution becomes basic enough that the indicator loses its hydrogen ion, flips to its deprotonated form, and shows the second color. The reverse happens when you titrate a base with an acid.
This is a chemical equilibrium. You can think of it like a seesaw: when hydrogen ion concentration is high, the balance tips toward the protonated color. When hydrogen ion concentration drops, it tips the other way. The color change isn’t instant across one drop of titrant. It happens over a narrow pH range, typically spanning about 1.5 to 2 pH units.
Why the Color Shifts at a Specific pH
Every indicator has a characteristic pH range where its color transition occurs. This range depends on how tightly the molecule holds onto its hydrogen ion, described by the indicator’s own dissociation constant. Methyl orange transitions between pH 3.1 and 4.4. Phenolphthalein transitions between pH 8.2 and 10.0. Litmus shifts around pH 5 to 8.
Within the transition range, both forms of the indicator coexist. You see an intermediate or blended color. Below the range, one form dominates completely. Above it, the other form dominates. The human eye generally detects the color change when the ratio of the two forms reaches roughly 1 to 10 in either direction. That’s why the visible shift covers a narrow window rather than a single exact pH value.
What Happens at the Equivalence Point
The equivalence point of a titration is where the acid and base have completely reacted with each other in stoichiometric proportions. At this moment, the pH of the solution changes very rapidly with each tiny addition of titrant. A single drop might swing the pH by several units. This dramatic jump is what makes the indicator appear to change color “suddenly” rather than gradually.
If you were to plot pH against volume of titrant added, you’d see a steep, nearly vertical section of the curve right around the equivalence point. The indicator’s transition range sits within that steep section, so the solution blows through the entire color-change window in just a drop or two. That’s what creates the sharp, satisfying endpoint you watch for in the lab.
Before and after this steep region, pH changes slowly with each added drop, so the indicator stays locked in one color form. The rapid pH swing at equivalence is the key reason the color change appears so abrupt.
Choosing the Right Indicator
For the indicator to accurately mark the equivalence point, its transition range needs to fall within that steep part of the titration curve. Different types of titrations produce different equivalence point pH values, so you can’t use the same indicator for every titration.
- Strong acid + strong base: The equivalence point falls at pH 7. The steep portion of the curve is wide, spanning roughly pH 4 to 10, so many indicators work. Phenolphthalein and methyl orange are both common choices.
- Weak acid + strong base: The equivalence point is above pH 7 (often around 8 to 9) because the conjugate base of the weak acid makes the solution slightly basic. Phenolphthalein, which transitions around pH 8.2 to 10, is a good match here.
- Strong acid + weak base: The equivalence point is below pH 7 (often around 4 to 6) because the conjugate acid of the weak base makes the solution slightly acidic. Methyl orange, transitioning around pH 3.1 to 4.4, is a better fit.
- Weak acid + weak base: The pH change near equivalence is gradual rather than steep, which means no indicator gives a sharp, reliable color change. These titrations are typically monitored with a pH meter instead.
If you pick an indicator whose transition range falls outside the steep section of the curve, the color will change too early or too late, giving you an inaccurate endpoint.
Why Only a Tiny Amount of Indicator Is Used
Since indicators are themselves weak acids or bases, adding too much would actually consume some of the titrant and throw off your results. You only need enough to produce a visible color. A few drops in a flask of solution is standard. At that tiny concentration, the indicator reacts to the pH of the solution around it without meaningfully affecting the titration chemistry.
Common Indicators and Their Color Changes
Phenolphthalein is probably the most widely used indicator in introductory chemistry. It goes from colorless in acidic solution to vivid pink in basic solution. The transition is easy to spot against a white background, which is one reason it’s so popular for teaching.
Methyl orange shifts from red in acidic conditions to yellow in basic conditions, with an orange intermediate. It’s useful for titrations with an acidic equivalence point. Bromothymol blue transitions from yellow (acidic) to blue (basic) around pH 6 to 7.6, making it a decent choice for strong acid/strong base work.
Universal indicator is a blend of several indicators that produces a continuous rainbow of colors across a wide pH range. It’s useful for estimating pH generally, but it’s not ideal for precise titration endpoints because the color transitions are less sharp than those of a single indicator.
Why the Color Sometimes Fades Back
If you’re titrating an acid with a base and you see phenolphthalein flash pink, then fade back to colorless, you haven’t yet reached the endpoint. The pink appears briefly where the drop of base hits the solution, creating a locally basic environment. But when you swirl and mix, the remaining acid neutralizes that base, pulling the pH back down and returning the indicator to its protonated, colorless form.
The true endpoint is reached when the pink color persists throughout the solution after thorough mixing, typically for at least 30 seconds. This persistence means the pH of the entire solution has crossed into the indicator’s transition range and stayed there. If you overshoot and add too much base, the color becomes deep and permanent rather than the faint pink of a proper endpoint.