A titration curve shows how the pH of a solution changes as a titrant is added from a burette. It is one of the most useful graphs in acid-base chemistry because it reveals the strength of the acid or base, the equivalence point, and the best indicator to use. The curve is usually steep near the equivalence point, where a small added volume causes a large pH change.
Reading these curves helps chemists measure unknown concentrations accurately.
Understanding Titration Curves
The shape of a titration curve comes from which particles are left in the flask after each small addition. At the start of a strong acid titration, there is a large excess of hydrogen ions, so the pH is low. Adding sodium hydroxide removes hydrogen ions by forming water.
Early additions produce a modest pH rise because many hydrogen ions remain. Near the reaction endpoint, very few excess hydrogen ions remain.
One extra drop can then change the balance sharply. After that point, hydroxide ions from the added base are in excess, so the pH is controlled mainly by the titrant.
Weak acids produce a different pattern because they do not fully split into ions in water. Before neutralisation is complete, the flask contains both the weak acid and its conjugate base. This mixture is a buffer.
It resists pH change because added hydroxide ions react with the weak acid, while added hydrogen ions react with the conjugate base. The buffer region is flatter than the corresponding part of a strong acid curve.
At half-equivalence, equal amounts of weak acid and conjugate base are present. This gives a useful way to find the acid dissociation constant from a graph, since the pH at this point equals the pKa.
The pH at equivalence depends on the salts left after reaction, not only on the original acid and base. A strong acid reacting with a strong base leaves ions that have little effect on water at room temperature. A weak acid reacting with a strong base leaves the conjugate base of the weak acid.
That conjugate base takes hydrogen ions from water and makes the solution alkaline. For this reason, its equivalence point lies above pH seven.
A weak base titrated with a strong acid has the opposite result. Its equivalence point lies below pH seven because the conjugate acid makes the solution acidic.
Indicators work because their acid form and base form have different colours. Each indicator changes over a limited pH interval rather than at one exact pH. The chosen indicator must complete its colour change during the rapid part of the curve.
Phenolphthalein is often suitable for a weak acid with a strong base because its change occurs in the alkaline range. Methyl orange can suit a strong acid with a weak base because it changes in the acidic range. Students should distinguish the endpoint from the equivalence point.
The endpoint is the colour change seen in the experiment. The equivalence point is the calculated chemical condition. Careful dropwise addition, swirling, and reading the burette at eye level help make those two points as close as possible.
Key Facts
- At the equivalence point, moles of acid and base have reacted in the exact stoichiometric ratio.
- For a strong acid and strong base titration, the equivalence point is at pH = 7 at 25 °C.
- Moles of solute = M x V, where M is molarity and V is volume in liters.
- For a monoprotic acid titrated with a strong base, MaVa = MbVb at the equivalence point.
- At the half-equivalence point of a weak acid titration, pH = pKa.
- A good indicator changes color within the steep vertical region near the equivalence point.
Vocabulary
- Titration
- A lab method used to determine an unknown concentration by reacting it with a solution of known concentration.
- Titrant
- The solution of known concentration that is added from the burette during a titration.
- Analyte
- The solution being tested, usually placed in the flask before the titration begins.
- Equivalence point
- The point where the amount of titrant added has completely reacted with the analyte according to the balanced equation.
- Buffer region
- A part of a weak acid or weak base titration curve where pH changes slowly because both a weak species and its conjugate are present.
Common Mistakes to Avoid
- Confusing the endpoint with the equivalence point: the endpoint is when the indicator changes color, while the equivalence point is the stoichiometric completion of the reaction.
- Using milliliters directly in mole calculations: volume must be converted to liters when using moles = M x V.
- Assuming every equivalence point has pH = 7: weak acid and weak base titrations have equivalence points above or below 7 because conjugate ions affect pH.
- Choosing an indicator only by color: the indicator must change color over the steep pH range near the equivalence point to give an accurate result.
Practice Questions
- 1 A 25.0 mL sample of HCl is titrated with 0.100 M NaOH. The equivalence point occurs at 32.5 mL of NaOH. What is the molarity of the HCl?
- 2 A weak acid is titrated with NaOH. The half-equivalence point occurs after 18.0 mL of NaOH has been added, and the pH is 4.76. What is the pKa of the acid, and at what volume is the equivalence point?
- 3 A titration curve starts at low pH, has a buffer region, and reaches an equivalence point above pH 7. Identify the type of titration and explain why the equivalence point is basic.