Acid–Base Indicators: How They Change Colour

Pour a few drops of a clear liquid into a clear solution and it turns bright pink. Add a little acid and the pink vanishes. It looks like a magic trick, and for centuries it was treated almost as one. Today we know exactly what’s going on: a molecule is gaining or losing a single proton, and that tiny change rearranges its electrons enough to change the colour of light it absorbs.

That molecule is an indicator, and understanding how it works connects acids and bases to the chemistry of colour.

What an indicator is

An acid–base indicator is a substance that has different colours in acidic and alkaline solutions. It “indicates” whether a solution is acidic, neutral or alkaline, and many indicators tell you roughly how acidic or alkaline.

Almost all indicators are weak acids (or occasionally weak bases). Chemists write a general indicator as HIn, where In stands for the rest of the molecule.

The chemistry of the colour change

In water, an indicator sets up an equilibrium:

HIn ⇌ H⁺ + In⁻

The two forms, HIn (the acid form) and In⁻ (the conjugate base form), have different colours.

• In acid, there are lots of H⁺ ions. By Le Chatelier’s principle, the equilibrium shifts to the left, so the HIn colour dominates.
• In alkali, OH⁻ ions remove H⁺ (forming water). The equilibrium shifts to the right, so the In⁻ colour dominates.

For litmus, HIn is red and In⁻ is blue. So litmus is red in acid and blue in alkali.

Why losing a proton changes colour

Colour comes from which wavelengths of visible light a molecule absorbs. Many indicator molecules contain long chains of alternating single and double bonds, called conjugated systems, where electrons are spread over many atoms. The longer and more extended the conjugated system, the longer the wavelength of light absorbed.

When an indicator gains or loses a proton, the arrangement of those bonds changes. The conjugated system might get longer, shorter, or rearrange around a ring. That shifts the wavelength of light absorbed, and so the colour we see changes.

Phenolphthalein is a neat example. In acid, its central carbon atom is bonded to four groups and breaks the conjugation into separate rings. The molecule absorbs only ultraviolet light, so it looks colourless. In alkali, it loses protons and the central carbon becomes part of one large conjugated system spread over three rings. Now it absorbs green light, so it looks pink-magenta.

The colour-change range

An indicator doesn’t flip at one exact pH. It changes gradually over a range, usually about two pH units, centred on the indicator’s pKa (often called pKin).

When pH = pKin, the two forms are present in equal amounts and you see a mixture of the two colours. Your eye starts to notice one colour dominating when it outnumbers the other by about ten to one, which happens roughly one pH unit either side.

Indicator	Colour in acid	Colour in alkali	Change range (pH)
Methyl orange	red	yellow	3.1–4.4
Methyl red	red	yellow	4.4–6.2
Litmus	red	blue	about 4.5–8.3
Bromothymol blue	yellow	blue	6.0–7.6
Phenolphthalein	colourless	pink	8.2–10.0

That range is why indicators are chosen carefully for titrations. See how to choose an indicator.

Types of indicator

Single indicators

One dye, one colour change. Litmus, methyl orange and phenolphthalein are single indicators. They tell you whether a solution is above or below a certain pH, and they’re ideal for titrations where you need one sharp change.

Universal indicator

A mixture of several dyes chosen so that their colour changes overlap, giving a smooth rainbow of colours across the whole pH range: red in strong acid, through orange, yellow and green at neutral, to blue and violet in strong alkali. It’s excellent for estimating pH but useless for titrations. See universal indicator.

Indicator papers

Paper strips soaked in indicator and dried. Litmus paper comes in red and blue versions: blue litmus turns red in acid; red litmus turns blue in alkali. Universal indicator paper gives a colour you match to a printed chart.

Natural indicators

Many plants contain pigments that change colour with pH, especially anthocyanins, the red, purple and blue pigments of berries, red cabbage and flowers. Natural indicators include:

• Red cabbage juice: red in acid, purple near neutral, green-yellow in alkali. Try making it in red cabbage indicator.
• Turmeric: yellow in acid and neutral solutions, turning red-brown in alkali. A curry stain turning red when you wash it with soap is turmeric acting as an indicator.
• Beetroot, blueberries, black tea, grape juice: all change colour noticeably with pH.
• Hydrangea flowers: the flowers themselves change colour depending on soil pH and aluminium. In acidic soil, aluminium becomes available and the flowers turn blue; in alkaline soil they’re pink.

Litmus itself is a natural product, extracted from certain lichens. People have used lichen dyes for centuries; Robert Boyle described their colour changes with acids in the 1660s.

Where indicators are used

• Titrations: finding the exact point where an acid has neutralised a base. See acid–base titration.
• Pool and aquarium testing: phenol red test kits check that water is near pH 7.2–7.8.
• Soil testing: gardeners use indicator-based kits to decide whether to add lime.
• Medical tests: urine test strips use indicators to check pH, and some wound dressings change colour when infection changes the wound’s pH.
• Food freshness: experimental “smart” packaging uses natural indicators that change colour as spoiling food releases basic amines or acidic compounds.
• Education: indicators are the simplest way to make acid–base chemistry visible.

Indicators vs pH meters

	Indicator	pH meter
Precision	about ±0.5 pH (universal)	±0.01 pH or better
Cost	very low	higher, needs calibration
Coloured or cloudy samples	hard to read	works fine
Speed	instant	seconds
Titration end points	very good with the right indicator	excellent, and records the whole curve

For quick checks, indicators win. For accurate measurements, a pH meter is better.

A quick experiment

You can see an indicator equilibrium shift with your own eyes:

1. Add a few drops of phenolphthalein to a dilute sodium hydroxide solution. It turns pink.
2. Add dilute hydrochloric acid drop by drop while swirling. At the point of neutralisation, the pink disappears.
3. Add a drop more sodium hydroxide. The pink returns.

Each colour flip is the indicator equilibrium HIn ⇌ H⁺ + In⁻ being pushed one way, then the other. Wear eye protection throughout.

Common misconceptions

• “The indicator reacts with the acid and is used up.” It’s an equilibrium; the indicator switches back and forth reversibly.
• “Indicators change at pH 7.” Most don’t. Methyl orange changes around pH 4 and phenolphthalein around pH 9.
• “Colourless means neutral.” Phenolphthalein is colourless in neutral and acidic solutions.
• “More indicator gives a better result.” A couple of drops is enough; too much can affect a titration.

Key takeaways

• Indicators are weak acids whose acid form (HIn) and conjugate base (In⁻) have different colours.
• Adding acid or alkali shifts the equilibrium HIn ⇌ H⁺ + In⁻ and changes the colour we see.
• Each indicator changes over about two pH units, centred on its pKin.
• Natural pigments such as anthocyanins are indicators too.
• For the underlying pH ideas, see the pH scale explained.