One drop flips the colour: how a titration catches an unknown concentration

How to use this simulation

1. The flask holds 25 mL of acid and an indicator, the burette 0.10 M NaOH. Open the tap and the yellow dot on the right-hand graph walks along the pH curve.
2. Near the end point the message above turns to “go drop by drop”. Close the tap and press “Add one drop” to catch the exact moment the colour changes.
3. When the colour flips, “End-point reading” and “Calculated concentration” light up. Compare with “True concentration” to see your error.
4. Now change the indicator, then switch the acid to H₂SO₄. With a 1 : 2 mole ratio the same acid concentration needs twice as much alkali. Check it.

Filling a glass until it overflows

Say you do not know how much water a glass holds. Pour in a spoonful at a time, count the spoonfuls, and the one that makes it overflow tells you the size. The glass is the unknown, the spoon is the known measure, and the overflow is the signal.

A titration does exactly this in chemistry. The acid in the flask is the unknown, the alkali in the burette has a known concentration, and the “overflow” signal is the indicator changing colour. Read how many mL of alkali it took and the acid’s concentration follows.

In the simulation the flask always holds 25 mL of acid. Move the concentration slider and the equivalence point shifts away from 20.0 mL, yet the same technique catches it every time.

The words of titration

Mixing these terms up costs marks, so settle them once.

From burette to flask, step by step

In a real laboratory the steps go in a fixed order to get a trustworthy result.

• Rinse the burette with a little of the alkali, so leftover water does not dilute it.
• Fill it and check that no air bubble sits in the tip below the tap. A bubble ruins the reading.
• Pipette exactly 25 mL of acid into the flask and add a few drops of indicator. A white tile under the flask makes the colour change easy to see.
• Record the starting burette reading. Keep your eye level with the meniscus and read the bottom of the curve.
• Open the tap and swirl the flask as you add alkali. Slow down once the colour starts to linger.
• Near the end point add single drops (one drop is about 0.05 mL). Stop at the first colour that lasts and record the final reading.
• Final minus starting reading is the titre. Repeat until you have concordant results, within 0.10 mL, and average them.

Indicators: why the colour changes

An indicator is itself a weak acid or base whose molecule and ion have different colours. As the pH moves, the ratio of the two forms moves, and so does the colour. Each indicator has a short pH range in which this happens.

Phenolphthalein changes between pH 8.2 and 10.0: colourless below, pink above. Methyl orange changes between 3.1 and 4.4: red below, yellow above, orange in between. The coloured bands on the graph in the simulation are exactly these ranges.

The pH curve: a leap in one drop

Titrate 25 mL of 0.10 M HCl with 0.10 M NaOH and the equivalence point sits at 25.0 mL. For almost the whole run the pH creeps upwards, then near that point it shoots almost straight up.

How big is the jump? One drop (0.05 mL) before the equivalence point the pH is 4.00; one drop after, it is 10.00. That is about seven pH units across a single drop, and it is why one drop flips the indicator completely.

The concentration formula

Molarity is moles per litre. Volume times molarity (in mL, millimoles) gives the amount. At the equivalence point the moles of H⁺ from the acid equal the moles of OH⁻ from the base, and that gives the rule.

For a 1 : 1 reaction like HCl and NaOH it is simply V₁M₁ = V₂M₂. But each H₂SO₄ molecule gives two H⁺, so one mole of acid needs two moles of NaOH, and the mole ratio has to go in.

Try it yourself in the simulation

Run these one at a time, and write down a prediction before each one.

• Experiment 1: HCl, phenolphthalein, 0.08 M. Open the tap. Near 20.0 mL the flask turns pink, and the calculated concentration lands close to the true one.
• Experiment 2: pour the whole lot with the tap fully open, then repeat going drop by drop near the end point. Which run has the smaller error?
• Experiment 3: switch to methyl orange. It turns from red to orange-yellow just before the equivalence point, so the calculated concentration comes out slightly low.
• Experiment 4: switch to H₂SO₄ at the same 0.08 M. The equivalence point is now 40.0 mL, exactly double. That is the 1 : 2 ratio.
• Experiment 5: raise the concentration until the end point would lie beyond the burette. The tap shuts at the end of the scale, as a real burette would run out.

Solved problems

Every number here is computed from the simulation’s own model, so you can check it against the screen.

Mistakes everyone makes

These are the places that cost the most marks, in exams and in the lab.

• Forgetting the mole ratio: using V₁M₁ = V₂M₂ directly for H₂SO₄ or H₃PO₄.
• Stopping after the end point: a deep pink means extra alkali and a concentration that comes out too high.
• Leaving a bubble in the burette tip, which makes the reading wrong.
• Not rinsing the pipette or burette: the water inside changes the concentration.
• Reading the meniscus from above: read the bottom of the curve with your eye level.
• Doing the titration once: you need at least three close readings.
• Choosing the wrong indicator: with a weak acid and a weak base no indicator gives a sharp end point.

Titration in real life

The method has long left the school lab and reached factories, hospitals and food shops.

• Medicines: tablet makers use titration to check how much active ingredient each batch holds.
• Food and drink: the acetic acid percentage in vinegar, the acidity of juice and how sour milk has become.
• Water testing: the alkalinity and hardness of river or pond water.
• Farming: measuring soil acidity to decide how much lime to add.
• Health: working out how much stomach acid an antacid tablet can neutralise.

Exam corner (SSC, HSC)

Try these yourself first, then check the answers.

• Short questions: What is a titration? How do the end point and equivalence point differ? Why must the burette and pipette be rinsed?
• Numerical: 25 mL of HCl needed 20.40 mL of 0.10 M NaOH. Find the molarity. Answer: 0.0816 M.
• Numerical: 25 mL of H₂SO₄ needed 18.50 mL of 0.10 M NaOH. Answer: 0.0370 M. Did you include the ratio?
• Reasoning: why is phenolphthalein good for a weak acid against a strong base? The solution at equivalence is basic, pH above 7, where its band lies.
• Sketch: draw the pH curve for a strong acid and a strong base and mark the equivalence point at pH 7.

Revision: at a glance

Read this list the night before the exam.

• A titration finds an unknown concentration (analyte) using a solution of known concentration (titrant).
• The end point is the indicator’s colour change; the equivalence point is the theoretical point of equal moles.
• For a 1 : 1 reaction, V₁M₁ = V₂M₂; otherwise bring in the mole ratio.
• For a strong acid and a strong base the equivalence pH is 7, and the curve is almost vertical there.
• The indicator’s pH range must lie inside the steep part of the curve.
• Go drop by drop at the end, read the bottom of the meniscus, and repeat three times.

Frequently asked questions