Acid–base indicators are essential tools in IB Chemistry Topic 8 (Acids and Bases) and in all titration work. They allow you to visualize pH changes through clear color shifts. Although indicators seem simple in practice, their behavior is grounded in equilibrium chemistry. This article explains exactly how indicators work, including their chemical structure, equilibrium shifts, and pH transition ranges.
What Is an Acid–Base Indicator?
An acid–base indicator is a weak acid or weak base that changes color depending on whether it is in its protonated or deprotonated form.
Most indicators have two forms:
- HIn (acid form)
- In⁻ (base form)
These two forms have different colors due to differences in molecular structure and electron distribution.
Examples:
- Phenolphthalein
- Methyl orange
- Bromothymol blue
- Litmus
Each indicator has its own color-changing range.
The Chemistry Behind Indicator Color Change
Indicators work because they establish an equilibrium:
HIn ⇌ H⁺ + In⁻
- HIn = protonated form (one color)
- In⁻ = deprotonated form (another color)
When pH changes, the position of this equilibrium shifts.
In acidic conditions (high [H⁺]):
Equilibrium shifts left:
- More HIn
- Color of the acid form appears
In basic conditions (low [H⁺]):
Equilibrium shifts right:
- More In⁻
- Color of the basic form appears
The visible color depends on which species predominates.
Why Indicators Have Distinct Colors
The color difference arises from structural changes in the indicator molecule.
- Protonated forms often have localized electrons
- Deprotonated forms often have more delocalization
- This changes how the molecule absorbs light
Small structural changes lead to large shifts in perceived color, making indicators effective and easy to observe.
Indicator Transition Range
Each indicator changes color over a narrow pH range, typically about 2 pH units.
This is called the transition range or working range, usually centered around the pKa of the indicator.
For an indicator HIn:
pH ≈ pKa → equal amounts of HIn and In⁻ (mixed color)
Color dominance:
- pH < pKa − 1 → mostly HIn (acid color)
- pH > pKa + 1 → mostly In⁻ (base color)
This explains why indicators do not change color instantly but gradually across a small interval.
Choosing the Right Indicator for a Titration
For a titration, the indicator’s transition range should match the equivalence point pH.
- Strong acid + strong base → equivalence near pH 7
Use indicators like bromothymol blue. - Strong acid + weak base → acidic equivalence
Use methyl orange. - Weak acid + strong base → basic equivalence
Use phenolphthalein. - Weak acid + weak base
No sharp pH jump → indicators do not work well; pH meter preferred.
IB exams often ask which indicator is most suitable and why.
Visual Example: Phenolphthalein
Phenolphthalein:
- Colorless in acidic solution (HIn)
- Pink in basic solution (In⁻)
Transition range: pH 8.2 to 10.0
This makes it ideal for strong base titrations.
Visual Example: Methyl Orange
Methyl orange:
- Red in acid
- Yellow in base
Transition range: pH 3.1 to 4.4
Suitable for strong acid titrations.
Indicators as Weak Acids or Weak Bases
Because indicators are weak acids or bases:
- They partially dissociate
- Their Ka (or Kb) determines their pKa and transition range
- They respond rapidly to pH changes
Understanding their equilibrium behavior helps explain their sensitivity.
Common IB Misconceptions
“Indicators measure pH directly.”
False. They only show whether the solution is above or below a certain range.
“Indicators work for all titrations.”
False. Weak acid–weak base titrations have no steep pH change.
“Indicators change at one exact pH value.”
They change over a range, not a single number.
FAQs
Why do indicators have two colors?
Their protonated and deprotonated forms absorb light differently, giving distinct colors.
Can indicators be used in non-aqueous solvents?
Usually no—most indicators rely on aqueous acid–base chemistry.
Why does the color change happen suddenly?
Although the shift is gradual, the eye perceives it as sudden because the equilibrium changes rapidly near pKa.
Conclusion
Indicators work because they are weak acids or bases whose protonated and deprotonated forms have different colors. Changes in pH shift the equilibrium between these forms, producing a visible color change. Understanding how indicators behave allows you to choose the correct one for titrations and interpret acid–base reactions confidently in IB Chemistry.
