neoCatalyst Labs · Chemistry · Free demo

Acid-Base Titration

Add base from a burette drop by drop into an acid solution. Watch the pH S-curve form in real time, identify the equivalence point, and calculate the unknown concentration.

Add base drop by drop and watch the pH curve form. Find the equivalence point in your own data.

6 acid-base pairs

~5 minutes

Lab report included

Learning Objectives

→ Add titrant drop-by-drop from a burette and identify the equivalence point from the pH S-curve
→ Calculate the unknown concentration of an acid using C₁V₁ = C₂V₂ stoichiometry
→ Run replicate trials and confirm concordance (agreement within 0.20 mL) before averaging your equivalence volume
→ Explain why equivalence point pH differs for strong acid/base vs. weak acid/base titrations

NGSS HS-PS1-7 AP Chemistry ⏱ 45 min

Student Name:

📖 Pre-Lab Reading — Acid-Base Titration

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What Is Titration?

Titration is a technique for determining the unknown concentration of a solution by reacting it with a precisely measured volume of a solution whose concentration is known. The solution in the burette (the titrant) is added drop by drop to the solution in the flask (the analyte) until the reaction is exactly complete.

The Neutralisation Reaction Acid + Base → Salt + Water. The equivalence point is the exact moment when moles of H⁺ equal moles of OH⁻ — neither acid nor base is in excess. Everything has been converted to salt and water.

Why the pH Changes the Way It Does

This is the key question to understand. Think about what is happening to the concentration of remaining acid as you add base:

Early in the titration (far from equivalence): You have a large amount of acid, so adding 1 mL of base removes only a tiny fraction of the H⁺ present. The concentration barely changes, so the pH barely changes. The curve is nearly flat.
Near the equivalence point: Almost all the acid has been consumed. The remaining moles of H⁺ are now very small. Adding one more drop of base destroys a large fraction of what's left — so the concentration (and therefore pH) changes dramatically with each tiny addition. This is why the curve goes nearly vertical.
After the equivalence point: All acid is gone. Further base just adds excess OH⁻, and you are in a new, large-volume solution. Each added mL changes the OH⁻ concentration only slightly. The curve flattens out again near pH 13–14.

The short answer: The dramatic jump near the equivalence point happens because a tiny addition of base destroys the last remaining acid — switching the solution from a tiny excess of H⁺ to a tiny excess of OH⁻. That single drop can shift pH from ~4 to ~10.

Why the Equivalence Point pH Depends on the Acid Type

At the equivalence point, all the acid has reacted with all the base. But what is left in the flask is not just water — it is a salt dissolved in water. Whether that salt makes the solution acidic, neutral, or basic determines the equivalence pH.

Strong acid + Strong base (e.g. HCl + NaOH): The salt formed is NaCl. Neither Na⁺ nor Cl⁻ reacts with water to produce H⁺ or OH⁻. The solution is neutral — pH = 7.
Weak acid + Strong base (e.g. acetic acid + NaOH): The salt formed contains the conjugate base of the weak acid (e.g. CH₃COO⁻, acetate). This conjugate base is itself weakly basic — it reacts with water: CH₃COO⁻ + H₂O → CH₃COOH + OH⁻. This produces a small amount of OH⁻, making the solution basic — pH > 7 at equivalence. The weaker the acid (larger pKₐ), the more basic the equivalence point.

Remember: A weak acid's conjugate base is not neutral — it is weakly basic. So when all the weak acid has been converted to its conjugate base by the strong base, the resulting solution is basic, not neutral.

Indicators and Endpoint

A pH indicator changes colour at a specific pH range. Phenolphthalein is colourless below pH 8.2 and pink above pH 10. Because the equivalence zone (pH 7–9 for most of our titrations) falls within this range and the pH jump is very steep, a single drop carries the indicator past its colour-change point. For weak acid + strong base, the equivalence pH is above 7, so phenolphthalein is still appropriate.

The endpoint is when the indicator first permanently changes colour. The equivalence point is the exact theoretical neutralisation point. A well-executed titration makes these nearly identical.

Indicator	pH Range	Colour Change	Best For
Phenolphthalein	8.2 – 10.0	Colourless → Pink	Strong acid + Strong base; Weak acid + Strong base
Methyl orange	3.1 – 4.4	Red → Orange/Yellow	Strong acid + Weak base
Litmus	6.0 – 8.0	Red → Blue	General use

Calculating Concentration from Titration

Once you know the equivalence volume, use the relationship that moles of acid = moles of base at equivalence:

C_acid × V_acid = C_base × V_base(eq) → C_acid = (C_base × V_base(eq)) / V_acid

Example: If 0.100 M NaOH required 24.8 mL to neutralise 25.0 mL of HCl, then C_HCl = (0.100 × 24.8) / 25.0 = 0.0992 M. This only works for monoprotic acids (one H⁺ per molecule).

Setup Today's Titration

Step 1 Perform the Titration

Add base from the burette in small increments. Watch the pH meter and the curve forming on the graph. When you see a large pH jump, you are near the equivalence point — switch to smaller additions.

Use these buttons to add base.

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Acidic ← → Basic

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Titration Curve — pH vs Volume of Base Added

Add Base: Added: 0.0 mL | Remaining: —

Add base from the burette. Start with 2 mL or 5 mL increments until the pH begins to rise quickly.

Step 2 Analysis Questions

Question 1

At what volume of base did the colour change occur (the endpoint)? How does this compare to your recorded equivalence point volume?

Mark your equivalence point in the data table first.

Question 2

What was the pH at the equivalence point? Was it 7, above 7, or below 7? Explain why based on the type of acid and base used.

Mark your equivalence point in the data table first.

Question 3 — Concentration Calculation

Mark equivalence point to see calculation.

Use C₁V₁ = C₂V₂ to calculate the concentration of the acid. Show your work.

Question 4

Why does the pH change slowly at the beginning of the titration, then rise dramatically near the equivalence point, then level off again?

Question 5

Conclusion

Write a conclusion synthesising your findings. Address: (1) what the shape of the pH S-curve reveals about acid-base reactions, (2) how the equivalence point pH relates to whether a strong or weak acid was used, and (3) what sources of error might affect your calculated concentration and how you would minimise them in a real lab.

Challenge Questions

Challenge 1

You performed a titration where 24.8 mL of 0.100 M NaOH neutralised 25.0 mL of HCl. Now suppose you repeated the experiment but accidentally diluted the NaOH solution to 0.090 M before starting. Without recalculating from scratch, predict whether the equivalence point would be reached at a smaller or larger volume of base, and by approximately what factor. Explain your reasoning using the relationship C₁V₁ = C₂V₂.

Challenge 2

A student performs a strong acid/strong base titration and another student performs a weak acid/strong base titration, both using the same concentrations and volumes. They compare their titration curves. Describe at least three specific differences you would expect to see between the two curves — not just the equivalence point pH, but also the shape of the curve before and after equivalence. Explain the chemistry behind each difference.

Challenge 3

Phenolphthalein changes colour between pH 8.2 and 10.0. For a weak acid/strong base titration where the equivalence pH is 8.7, explain why phenolphthalein is an appropriate indicator and what error (if any) it introduces. Then name one scenario — a different acid-base system — where phenolphthalein would be an inappropriate indicator choice, and recommend a better alternative. Justify your answer using the concept of the indicator colour-change range vs. the equivalence pH.

Run at least two titration trials and answer all 5 analysis questions to enable submission.

Course Navigator — Unit K: Acids, Bases & pH

✓ K-1 pH Scale → ✓ K-2 Red Cabbage → ✦ K-3 Titration

This lab completes Unit K. You've progressed from measuring pH qualitatively (K-1), to observing natural indicators (K-2), to the quantitative technique of titration (K-3). Titration forms the foundation of analytical chemistry — the next unit builds on these stoichiometry skills.

Acid-Base Titration

📖 Pre-Lab Reading — Acid-Base Titration

What Is Titration?

Why the pH Changes the Way It Does

Why the Equivalence Point pH Depends on the Acid Type

Indicators and Endpoint

Calculating Concentration from Titration

What your student just did

What you see as the parent

Marking guide

Generate the lab report

Where that mark ends up