Last updated: 18 August 2026
In an acid–base titration, the exact stoichiometric point is invisible. The acid does not suddenly disappear, and the base does not announce that exactly enough has been added. We therefore need some way of detecting when the titration has reached the region very close to equivalence.
A pH indicator provides that visible signal by changing colour as the pH changes. However, an indicator does not switch from one colour to another at one exact pH. It changes gradually over a range of pH values.
Choosing a suitable indicator is therefore a matching problem. We must understand the equilibrium responsible for the indicator colour change, determine its transition range, and compare that range with the steep portion of the titration curve.
Try these five questions before reading the chapter. They check the acid–base ideas that make indicator choice easier to understand.
- 5 short questions
- Instant explanation for every option
- No marks are lost — this is a starting-point check
How acid–base indicators work
An acid–base indicator is normally a weak organic acid or weak organic base whose protonated and deprotonated forms have different colours.
We can represent a weak acid indicator as HIn. It partially ionises in water:
The symbol In represents the rest of the indicator molecule. The important point is that HIn and In⁻ have different electronic structures and therefore absorb visible light differently. As a result, they appear as different colours.
| Indicator form | Favoured conditions | Observed colour |
|---|---|---|
| HIn | Lower pH, higher [H⁺] | Colour of the acidic form |
| In⁻ | Higher pH, lower [H⁺] | Colour of the alkaline form |
What happens when the solution becomes more acidic?
Adding acid increases [H⁺]. For the equilibrium
the increase in [H⁺] favours the reverse reaction. The proportion of HIn increases, so the colour associated with HIn becomes dominant.
What happens when the solution becomes more alkaline?
In alkaline solution, OH⁻ removes H⁺:
As [H⁺] decreases, the indicator equilibrium shifts towards the right. More In⁻ is produced and the colour associated with In⁻ becomes dominant.
HIn and In⁻ coexist in equilibrium. Changing the pH changes their relative concentrations. The visible colour changes because one form gradually becomes much more abundant than the other.
The indicator constant and pKIn
Since HIn behaves as a weak acid, we can write an equilibrium constant for its ionisation:
Rearranging:
Taking negative logarithms gives:
This is simply the Henderson–Hasselbalch relationship applied to the indicator. It connects the pH directly to the ratio of the two coloured forms.
What happens when pH = pKIn?
Suppose:
Then:
and:
Therefore:
At this point, the acidic and alkaline forms are present in equal concentrations. The colour seen is normally intermediate between the two limiting colours.
It is not a magical pH at which the indicator suddenly changes colour. The transition takes place over a range centred approximately on pKIn.
Why an indicator changes colour over a pH range
Our eyes do not detect a clear colour change as soon as a tiny amount of the second form appears. One form generally needs to become much more abundant than the other before its colour dominates.
A useful approximation is that the acidic colour dominates when:
Using:
gives:
At the other end of the visible transition, the alkaline form dominates when:
giving:
Hence the approximate working range of an indicator is:
| pH | [In⁻] : [HIn] | Main form present | Colour seen |
|---|---|---|---|
| pKIn − 1 | 1 : 10 | Mostly HIn | Acidic colour dominates |
| pKIn | 1 : 1 | Equal proportions | Intermediate colour |
| pKIn + 1 | 10 : 1 | Mostly In⁻ | Alkaline colour dominates |
Incorrect idea: An indicator changes from one colour to another exactly when pH = pKIn.
Correct reasoning: pKIn is approximately the midpoint of the colour change. The visible transition normally extends over roughly pKIn ± 1.
Equivalence point and end point
Two ideas that sound similar need to be separated carefully: equivalence point and end point.
The equivalence point
The equivalence point is defined by the stoichiometry of the reaction. At equivalence, the acid and base have reacted in exactly the mole ratio required by the balanced equation.
For example, for:
equivalence is reached when the reacting amounts of H⁺ and OH⁻ are equal.
Notice that this definition says nothing about the colour of an indicator.
The end point
The end point is the experimentally observed colour change used to decide when to stop adding titrant.
Since an indicator changes colour over a range of pH values, its end point does not necessarily occur at exactly the same pH as the theoretical equivalence point.
This happens when the indicator changes colour within the steep portion of the titration curve, where a very small change in titrant volume causes a very large change in pH.
Choosing a suitable indicator from a titration curve
The most reliable method is not to memorise isolated rules such as “methyl orange for acids” or “phenolphthalein for bases”. Instead, look at the shape of the titration curve.
Near equivalence, a suitable acid–base titration shows a rapid change in pH. On a graph of pH against volume of titrant added, this appears as a steep, nearly vertical section.
A reliable four-step method
- Locate the steep section. This is the rapid pH change around equivalence.
- Read its approximate pH limits. Determine the pH interval covered by the steep portion.
- Check the indicator range. Use the range supplied in the question or estimate it from pKIn ± 1.
- Match the ranges. Choose an indicator whose transition occurs within the steep portion of the titration curve.
What matters is that its transition takes place while the titration curve is very steep. Under those conditions, the whole colour change corresponds to only a very small change in titrant volume.
Methyl orange and phenolphthalein
Questions may provide the transition range of an unfamiliar indicator. If so, use the data given rather than relying on memory. Two familiar examples, however, are methyl orange and phenolphthalein.
| Indicator | Approximate pH range | More acidic colour | More alkaline colour |
|---|---|---|---|
| Methyl orange | 3.1–4.4 | Red | Yellow |
| Phenolphthalein | 8.2–10.0 | Colourless | Pink |
Methyl orange therefore changes colour in an acidic pH region, whereas phenolphthalein changes colour in an alkaline pH region.
Strong acid–strong base titrations
In a strong acid–strong base titration, both substances are essentially fully ionised. Before equivalence there is excess strong acid; after equivalence there is excess strong base.
Near equivalence, therefore, a tiny addition of titrant can change the solution rapidly from distinctly acidic to distinctly alkaline. The pH jump is very large.
For a strong monoprotic acid titrated with a strong monobasic base at 25°C, the equivalence point is at approximately pH 7.
However, this does not mean that the indicator must change at pH 7. The steep region may extend over many pH units.
Methyl orange: its transition range can lie within the lower part of the steep pH jump.
Phenolphthalein: its transition range can lie within the upper part of the same steep pH jump.
Therefore, both may be suitable for a strong acid–strong base titration.
Strong acid–weak base titrations
Consider a weak base, B, reacting with a strong acid:
At equivalence, the weak base has largely been converted into its conjugate acid, BH⁺. The conjugate acid can react with water:
H₃O⁺ is produced, so the solution at equivalence is acidic. Its pH is therefore below 7.
The steep pH change occurs in a more acidic region than for a strong acid–strong base titration.
An indicator with an acidic transition range, such as methyl orange, is therefore often suitable. Phenolphthalein normally changes colour outside the useful steep region and is unsuitable.
The relevant question is where the steep pH change occurs. For a strong acid–weak base titration, that region is shifted towards lower pH.
Weak acid–strong base titrations
Now consider a weak acid, HA, reacting with a strong base:
At equivalence, most of the weak acid has been converted into its conjugate base, A⁻. This ion undergoes hydrolysis:
The production of OH⁻ means that the equivalence solution is alkaline, with pH above 7.
The steep pH change is therefore shifted towards higher pH. An indicator such as phenolphthalein is usually suitable.
Methyl orange changes colour at much lower pH and would normally reach its end point significantly before the equivalence region.
Weak acid–weak base titrations
When a weak acid is titrated with a weak base, neither reagent produces a large concentration of H⁺ or OH⁻. Consequently, the pH changes much more gradually near the equivalence point.
There is usually no sufficiently steep vertical section for an ordinary visual indicator to complete its transition over a negligible change in titrant volume.
An instrumental method such as monitoring the pH with a pH meter is preferable because the equivalence region cannot be located accurately from a conventional indicator colour change.
Comparing the four common titration types
| Titration | Equivalence region | Typical indicator choice | Reason |
|---|---|---|---|
| Strong acid + strong base | Around pH 7 at 25°C, with a very large pH jump | Methyl orange or phenolphthalein may both work | Both ranges can fall within the steep section |
| Strong acid + weak base | Below pH 7 | Methyl orange or another acidic-range indicator | The rapid pH change occurs at relatively low pH |
| Weak acid + strong base | Above pH 7 | Phenolphthalein or another alkaline-range indicator | The rapid pH change occurs at relatively high pH |
| Weak acid + weak base | No sufficiently sharp pH jump in most cases | No simple visual indicator | The pH changes too gradually near equivalence |
Worked examples: choosing an indicator
Example 1: a large pH jump
A titration curve changes almost vertically from approximately pH 3.0 to pH 11.0. Two available indicators are:
- methyl orange: pH 3.1–4.4;
- phenolphthalein: pH 8.2–10.0.
Both transition ranges lie within the rapid pH change.
Conclusion: Both indicators are suitable. There is no requirement for the transition range to be centred on pH 7.
Example 2: strong acid–weak base
A titration curve shows its rapid pH change mainly between pH 3 and pH 6.
Methyl orange changes between approximately pH 3.1 and 4.4, so its transition lies within the steep region.
Phenolphthalein changes between approximately pH 8.2 and 10.0, well outside the steep region.
Conclusion: Methyl orange is suitable; phenolphthalein is not.
Example 3: weak acid–strong base
Suppose the steep pH change occurs approximately from pH 7.5 to pH 11.
Phenolphthalein, with a transition range of approximately pH 8.2–10.0, changes colour within this steep region.
Methyl orange would have completed its colour change at a much lower pH, before equivalence is reached.
Conclusion: Phenolphthalein is suitable; methyl orange is not.
Using pKIn in calculations
Example 1: finding the approximate indicator range
An indicator has:
The approximate colour-change range is:
Answer: approximate transition range = pH 4.2–6.2.
Example 2: finding the ratio of indicator forms
An indicator has pKIn = 4.5. Find [In⁻] : [HIn] at pH 5.5.
Substituting:
Answer: [In⁻] : [HIn] = 10 : 1, so the alkaline form predominates.
Does reversing the titration change the indicator choice?
If an alkali is added to an acid, the titration curve rises. If acid is added to an alkali, the corresponding curve falls.
Reversing the direction changes the direction of the colour change, but it does not fundamentally change the pH region in which equivalence occurs.
Rising pH: the indicator changes from its lower-pH colour towards its higher-pH colour.
Falling pH: the same indicator changes in the opposite direction.
In both cases, suitability still depends on whether its transition range lies within the steep section of the titration curve.
Why an unsuitable indicator gives an inaccurate titre
Suppose the transition range of an indicator lies well outside the steep portion of the titration curve. A noticeable volume of titrant may then have to be added before the colour change is reached, or the colour change may occur noticeably before equivalence.
The recorded titre would therefore differ significantly from the true equivalence volume.
By contrast, when the indicator range lies in the steep region, a change of several pH units corresponds to only a tiny change in titrant volume.
The important feature is the steepness of the titration curve. This explains why a suitable indicator need not change colour exactly at the equivalence-point pH.
Common mistakes when choosing pH indicators
| Common mistake | Why it is wrong | Correct approach |
|---|---|---|
| Assuming every indicator changes colour at pH 7 | Different indicators have different KIn values | Use the stated range or pKIn ± 1 |
| Thinking the indicator changes instantaneously at pKIn | The HIn : In⁻ ratio changes continuously with pH | Treat pKIn as the midpoint of a transition range |
| Choosing the indicator whose pKIn is closest to the equivalence-point pH without considering the curve | The whole indicator range may still lie within the steep section even if its midpoint is not close to the equivalence pH | Match the transition range to the steep pH change |
| Using phenolphthalein automatically whenever a base is involved | A strong acid–weak base titration has an acidic equivalence region | Determine where the rapid pH change occurs |
| Using methyl orange automatically whenever an acid is involved | A weak acid–strong base titration has an alkaline equivalence region | Use the titration curve rather than the initial reagent names |
| Saying the end point and equivalence point are identical | The end point is observed; equivalence is defined stoichiometrically | State that a suitable indicator makes them correspond very closely in volume |
| Rejecting an indicator because it does not change colour exactly at the equivalence-point pH | An indicator may still change entirely within the steep pH jump | Judge the indicator against the steep region |
| Choosing a visual indicator for a weak acid–weak base titration simply from the equivalence pH | The pH change near equivalence is normally too gradual | Use an instrumental method such as a pH meter |
| Thinking reversing the burette and flask solutions automatically requires a different indicator | The curve direction reverses, but the relevant equivalence pH region remains similar | Match the same indicator range to the steep section |
pH-indicator checklist
Understand the equilibrium
- Write HIn ⇌ H⁺ + In⁻.
- Remember that HIn and In⁻ have different colours.
- Higher [H⁺] favours HIn.
- Lower [H⁺] favours In⁻.
Use pKIn
- At pH = pKIn, [HIn] = [In⁻].
- Use pH = pKIn + log([In⁻]/[HIn]).
- Estimate the transition range as pKIn ± 1.
- Do not treat the colour change as instantaneous.
Read the titration curve
- Locate the steep pH change around equivalence.
- Estimate the pH interval of this steep section.
- Compare it with the indicator range.
- Choose an indicator changing within the steep region.
Recognise the titration type
- Strong acid–strong base: several indicators may work.
- Strong acid–weak base: use an acidic-range indicator.
- Weak acid–strong base: use an alkaline-range indicator.
- Weak acid–weak base: a visual indicator is normally unsuitable.
Now apply the chapter to unfamiliar titration curves, indicator ranges and equilibrium calculations.
- 10 application-focused questions
- Feedback explains every answer choice
- Use your score to identify what to revise
Frequently asked questions
Questions students often ask about pH indicators
These questions connect indicator equilibria with the decisions you need to make when interpreting acid–base titration curves.
Does an indicator have to change colour exactly at the equivalence-point pH?
No. Its transition range needs to lie within the steep pH change around the equivalence point.
In that region, several pH units may correspond to only a tiny change in titrant volume. The indicator end point can therefore occur extremely close to the true equivalence volume even if the two pH values are not identical.
Why can phenolphthalein work in a strong acid–strong base titration if it changes above pH 7?
A strong acid–strong base titration has a very large pH jump. Although the equivalence point is around pH 7 at 25°C, the steep section can extend well into the alkaline region.
Phenolphthalein can therefore complete its transition while the curve is still extremely steep.
Why is the equivalence point below pH 7 for a strong acid–weak base titration?
At equivalence, the weak base has been converted largely into its conjugate acid. The conjugate acid reacts with water to produce H₃O⁺, making the solution acidic.
Why is the equivalence point above pH 7 for a weak acid–strong base titration?
At equivalence, the weak acid has been converted largely into its conjugate base. The conjugate base reacts with water to produce OH⁻, making the solution alkaline.
Why is the indicator range roughly two pH units wide?
A useful approximation is that one indicator colour dominates when one form is around ten times as concentrated as the other.
Changing [In⁻]/[HIn] from 0.1 to 10 changes log([In⁻]/[HIn]) from −1 to +1. The visible range therefore extends from about pKIn − 1 to pKIn + 1.
At pH = pKIn, has exactly half of the indicator ionised?
For the simple equilibrium HIn ⇌ H⁺ + In⁻, pH = pKIn means [HIn] = [In⁻]. The two indicator forms are therefore present in equal concentrations.
Is pKIn fundamentally different from pKa?
No. KIn is simply the acid dissociation constant used for the weak acid form of an indicator, HIn.
The notation pKIn helps remind us that the equilibrium belongs to an indicator.
What should I do if the exam gives an indicator I have never heard of?
You do not need to recognise its name. Use the transition range supplied in the question and compare it with the steep portion of the titration curve.
If only pKIn is supplied, estimate the range using pKIn ± 1.
Does reversing which solution is placed in the burette change the suitable indicator?
Not normally. Reversing the titration changes whether the curve rises or falls, but the steep equivalence region remains in essentially the same pH range for the same acid–base pair.
The direction of the indicator colour change simply reverses.
Why should only a small amount of indicator be used?
An acid–base indicator is itself an acid–base species. Only a small quantity is needed to produce a visible colour, so using just a few drops minimises any disturbance to the system being titrated.
Why can a weak acid–weak base titration be followed with a pH meter even when an indicator is unsuitable?
A pH meter measures pH continuously and does not depend on a visual colour transition occurring over a narrow volume interval.
The curve can therefore still be recorded even when the pH change near equivalence is too gradual for a conventional indicator to provide a sharp end point.
Choosing a pH indicator becomes much easier once you stop treating indicators as substances that suddenly change colour at one special pH. The two coloured forms exist in equilibrium, and changing pH changes their relative proportions. This gives each indicator a transition range of approximately pKIn ± 1. In a titration, the final step is to match that range to the steep pH change around equivalence. Once this idea is secure, indicator choice follows naturally from the chemistry and the shape of the titration curve rather than from memorised rules.