Phenolphthalein: Why It Turns Pink

Add a few drops of phenolphthalein to a colourless alkaline solution and a vivid pink-magenta blooms through the liquid. Add acid and it vanishes as if it had never been there. It’s the most recognisable colour change in school chemistry, and it’s the go-to indicator for titrating weak acids.

But phenolphthalein has more tricks than most students realise. It fades back to colourless in very strong alkali, it turns orange in concentrated sulfuric acid, and it was once sold as a laxative. All of these behaviours come from the same molecule changing its shape.

The basics

Property	Value
Formula	C₂₀H₁₄O₄
Molar mass	318.3 g/mol
Appearance (solid)	white to pale yellow crystals
Colour in acid and neutral solution	colourless
Colour in alkali	pink to magenta
Colour-change range	about pH 8.2–10.0
pKa (for the colour change)	about 9.4
Solubility	poorly soluble in water; used dissolved in ethanol

Because it barely dissolves in water, phenolphthalein indicator is sold as a solution in ethanol, typically about 0.5–1% by mass. That’s why adding a lot of it to water can make the solution slightly cloudy.

The structure

Phenolphthalein has a central carbon atom joined to three ring systems: two phenol rings (benzene rings each carrying an –OH group) and a third ring that’s part of a five-membered lactone (a cyclic ester).

The key to the colour is what happens at that central carbon.

Why it’s colourless in acid

In acidic and neutral solutions, the central carbon is sp³ hybridised: it has four single bonds, including one to the oxygen of the lactone ring. This tetrahedral carbon acts as an insulator. The three rings are each separate conjugated systems, and none of them is large enough to absorb visible light. The molecule only absorbs ultraviolet, so we see it as colourless.

Why it turns pink in alkali

As the pH rises above about 8.2, hydroxide ions remove the protons from the two phenol –OH groups. The resulting negative charge pushes electrons towards the central carbon, which breaks the lactone ring open. The central carbon becomes sp² hybridised and planar.

Now the three rings are joined into one large conjugated system, with electrons delocalised across all of them. That extended system absorbs light in the green part of the visible spectrum, around 550 nm. The light that passes through is mostly red and blue, which our eyes see as pink-magenta.

In simple equilibrium terms:

H₂In (colourless) ⇌ In²⁻ (pink) + 2H⁺

The deeper the pH goes into the 8.2–10.0 range, the more molecules are in the pink form, so the pink deepens from pale to intense.

Why the pink fades in very strong alkali

Leave a pink phenolphthalein solution in concentrated sodium hydroxide (pH above about 13) and the colour slowly fades back to colourless over minutes to hours.

This happens because hydroxide ions attack the central carbon directly and add to it, forming a new species, often written In(OH)³⁻. The central carbon becomes sp³ again, the big conjugated system is broken up, and the colour disappears. This reaction is slow, which is why the fading takes time.

This is a useful reminder that “colourless” with phenolphthalein doesn’t always mean acidic. At pH 14 it can be colourless too.

And in concentrated sulfuric acid?

In very strongly acidic conditions, such as concentrated sulfuric acid, phenolphthalein takes up an extra proton and forms a cation with a different conjugated system. It turns orange-red. You’ll rarely see this in school, but it shows the same principle: every change in how the molecule is protonated changes the conjugation, and every change in conjugation changes the colour.

Conditions	Form	Colour
Concentrated H₂SO₄	protonated cation	orange-red
pH 0 to about 8.2	lactone (H₂In)	colourless
pH 8.2 to about 12	open dianion (In²⁻)	pink to magenta
pH above about 13	hydroxide adduct (In(OH)³⁻)	slowly fades to colourless

Using phenolphthalein in titrations

Phenolphthalein’s colour change (8.2–10.0) sits exactly where the steep part of the titration curve is for a weak acid titrated with a strong base. At the equivalence point of ethanoic acid with sodium hydroxide, the pH is about 8.7, right in the middle of phenolphthalein’s range. See titration curves.

It also works for strong acid–strong base titrations, where the pH jump runs from about 3.5 to 10.5.

It is not suitable for titrating a weak base (such as ammonia) with a strong acid, because the equivalence point is around pH 5, well below its range. Use methyl orange there. See how to choose an indicator.

Practical tips:

• Titrate so that the colour change is colourless → pink (acid in the flask, alkali in the burette). Eyes detect the first appearance of pink more easily than its disappearance.
• Stop at the first permanent pale pink. A pink that fades after a few seconds usually means incomplete mixing, or carbon dioxide from the air slowly neutralising the last traces of alkali.
• Two or three drops is enough.

Other uses

Forensic science. The Kastle–Meyer test uses a reduced, colourless form of phenolphthalein (phenolphthalin) to detect traces of blood. Haemoglobin catalyses its oxidation by hydrogen peroxide back to phenolphthalein, which turns pink in the alkaline test solution. It’s a presumptive test: a positive result suggests blood, which is then confirmed by other methods.

Disappearing ink. A novelty ink made from phenolphthalein in a weakly alkaline solution looks pink when applied. As the ink dries, carbon dioxide from the air neutralises the alkali and the colour fades.

Concrete testing. Spraying phenolphthalein on freshly broken concrete shows how far carbonation has penetrated. Fresh concrete is strongly alkaline (pH about 12–13) and turns pink; carbonated concrete, which has reacted with CO₂ from the air, has a lower pH and stays colourless. Carbonation matters because it can allow steel reinforcement inside the concrete to corrode.

Former laxative. For most of the 20th century, phenolphthalein was the active ingredient in several over-the-counter laxatives. It was withdrawn in the United States and several other countries in the late 1990s after animal studies raised concerns about possible cancer risk. It’s now handled as a laboratory chemical only.

History

Phenolphthalein was first made in 1871 by the German chemist Adolf von Baeyer, who later won the 1905 Nobel Prize in Chemistry for his work on organic dyes. He made it by heating phthalic anhydride with phenol in the presence of an acid catalyst, a reaction that’s still the basis of its manufacture. The name comes from its two parent compounds: phenol and phthalic anhydride.

Safety

• The indicator solution is flammable because it’s mostly ethanol.
• Phenolphthalein itself is treated as a possible carcinogen and reproductive hazard, so avoid skin contact and never ingest it.
• In the tiny amounts used in titrations, the risk is very low with normal lab precautions: eye protection, gloves where required, and washing hands.

Key takeaways

• Phenolphthalein is colourless below about pH 8.2 and pink-magenta above it, up to about pH 12.
• The colour appears when the central carbon becomes sp², linking all three rings into one conjugated system that absorbs green light.
• In very strong alkali, hydroxide adds to the central carbon and the colour slowly fades.
• It’s the standard indicator for weak acid–strong base titrations.
• Compare it with other indicators in acid–base indicators.