Acid–base indicators make pH visible by undergoing a reversible acid–base reaction. Because the two conjugate forms have different colours (or other observable properties), the dominant form at a given pH determines what you see.

Core idea: two forms, two appearances

Most indicators are weak acids or weak bases that exist as an equilibrium mixture of two main species:

• a protonated form (more H attached)

• a deprotonated form (one fewer H, often carrying more negative charge)

Because these two forms have different electronic structures, they absorb and transmit different wavelengths of light, producing different colours.

A key consequence of the syllabus statement is that the indicator does not “measure pH directly”; instead, pH controls the indicator’s protonation state, and the protonation state controls the observed colour.

Indicator equilibrium (what pH is actually shifting)

For a common weak-acid indicator, the equilibrium can be represented as:

• HIn (protonated indicator) + H2O ⇌ In− (deprotonated indicator) + H3O+

The solution’s pH reflects the hydronium level, and the hydronium level pushes this equilibrium:

• Lower pH (more H3O+) shifts toward HIn (protonated form)

• Higher pH (less H3O+) shifts toward In− (deprotonated form)

This is why the same indicator can display different colours in acidic vs basic solutions: the dominant species changes as the equilibrium position changes.

Interpreting this expression qualitatively is the main skill: if pH changes, then [H3O+] changes, so the ratio [In−]/[HIn] must adjust to satisfy the equilibrium constant for that indicator.

Plot showing how an indicator’s spectroscopic signal varies systematically with pH for bromothymol blue, reflecting the changing proportions of $\mathrm{HIn}$ and $\mathrm{In^-}$. As pH increases, the equilibrium shifts toward the deprotonated form, producing a predictable, smooth change in measured absorbance that corresponds to the changing ratio $[In^-]/[HIn]$. Source

Why the colour change occurs over a range (not at one pH)

Indicators typically have a transition range because human-visible colour depends on having enough of one form to dominate the perceived colour.

Bromothymol blue in solutions of different pH, showing a gradual visible shift from the acid-form color (yellow) through intermediate blended colors (green) to the base-form color (blue). This provides a visual example of a transition range, where both $\mathrm{HIn}$ and $\mathrm{In^-}$ can be present in comparable amounts and the perceived color is a mixture. Source

• When one form is present in a small fraction, its colour is usually masked.

• When both forms are present in comparable amounts, the observed colour is often a blend.

A useful way to think about the transition:

• At pH values where the protonated form is much larger than the deprotonated form, you see the acid colour.

• At pH values where the deprotonated form is much larger, you see the base colour.

• In between, you see a mixed colour.

“Different properties” beyond colour

While colour is most common, the syllabus language is broader: protonated vs deprotonated forms can differ in other observable properties, such as:

• fluorescence (glowing under UV)

• absorbance spectrum (instrumental colour intensity)

• solubility/partitioning in some systems

In each case, the underlying reason is the same: protonation changes molecular structure and electron distribution, changing how the substance interacts with light or its environment.

Practical interpretation (what you should say in explanations)

When explaining indicator behaviour, focus on these cause-and-effect links:

• pH sets [H3O+]

• [H3O+] shifts the indicator equilibrium

• the dominant conjugate form determines the observed property (e.g., colour)

Use language like “predominantly protonated” or “predominantly deprotonated” rather than claiming the indicator is entirely one form, because it is an equilibrium mixture.