Titration is a laboratory technique used to determine the concentration of an unknown solution by reacting it with a solution of known concentration. It is important in chemistry because it connects precise measurement with balanced chemical equations. Students use titration to analyze acids, bases, oxidizing agents, metal ions, and many real samples such as vinegar or water.
Good technique matters because a small reading or mixing error can change the final calculated concentration.
In a volumetric titration, a burette delivers the standard solution, called the titrant, into a flask containing the analyte. The reaction is followed until the endpoint, often shown by a color change from an indicator or a reading from a pH meter. The volume delivered from the burette is combined with the balanced reaction ratio to calculate moles and concentration.
More advanced methods, such as back titration, are useful when the direct reaction is slow, incomplete, or difficult to detect at the endpoint.
Understanding Chemistry: Titration Techniques
Reliable results begin before any liquid enters the flask. The solution with a stated concentration must be trustworthy. Some chemicals absorb water or carbon dioxide from air, so their mass does not give a reliable number of moles.
Sodium hydroxide is a common example. Chemists often standardize such a solution against a primary standard. A primary standard is very pure, stable, and easy to weigh accurately.
Potassium hydrogen phthalate is often used for this purpose. Standardization finds the actual concentration of the working solution, which may differ slightly from the value written on its label. This step prevents every later calculation from being based on a hidden error.
Careful handling of glassware controls much of the uncertainty. A burette should be rinsed with small portions of the solution it will contain. Water left inside would dilute that solution.
The tip must be filled before recording the starting reading, since an air bubble can make the recorded volume larger than the amount that reached the flask. Read the bottom of the liquid curve at eye level. Looking from above or below causes parallax error.
Near the endpoint, add liquid one drop at a time while swirling continuously. Rinse drops from the flask wall with distilled water. This added water does not change the number of moles present, so it does not affect the result.
An indicator works because its color depends on chemical conditions in the flask. For acid base work, the best indicator changes color over the steep part of the pH change. Phenolphthalein is useful when its color range matches the reaction being studied.
A poor indicator can change before or after the true stoichiometric completion, creating a consistent error. The final faint color should persist for about thirty seconds, not become dark. A dark color usually means too much titrant was added.
A pH meter can give a more detailed record by showing how pH changes as volume is added. The sharpest part of this graph helps locate the equivalence volume.
Calculations must follow the balanced chemical equation rather than a memorized shortcut. First find the volume delivered from the difference between two burette readings. Use the known concentration and this volume to find moles of the known reactant.
Then use the coefficients in the equation to convert to moles of the unknown substance. Finally divide those moles by the sample volume in liters. Repeating the procedure gives concordant results when measured volumes are very close together.
Results that disagree may show random reading variation, incomplete mixing, contaminated equipment, or an overshot endpoint. These same skills matter in water testing, food analysis, medicine production, and checking the acidity of products such as vinegar.
Key Facts
- Molarity is concentration in moles per liter: M = n/V.
- For a reaction aA + bB -> products, mole ratio is nA/a = nB/b at equivalence.
- For many acid-base titrations with 1:1 stoichiometry, M1V1 = M2V2.
- Burette volume delivered equals final burette reading minus initial burette reading.
- The endpoint is the observed signal, while the equivalence point is the stoichiometric completion point.
- Percent error can be calculated as percent error = |experimental - accepted|/accepted x 100 percent.
Vocabulary
- Burette
- A graduated glass tube with a stopcock used to deliver a measured volume of solution accurately.
- Titrant
- The solution of known concentration added from the burette during a titration.
- Analyte
- The solution of unknown concentration being tested in the flask.
- Endpoint
- The point in a titration where an observable change shows that enough titrant has been added.
- Back titration
- A titration method in which excess known reagent is added to the sample and the leftover reagent is titrated.
Common Mistakes to Avoid
- Reading the meniscus from above or below eye level, which gives a parallax error and makes the delivered volume inaccurate.
- Forgetting to rinse the burette with the titrant, which can dilute the standard solution and lower the calculated concentration.
- Adding titrant too quickly near the endpoint, which can overshoot the color change and produce a volume that is too large.
- Using M1V1 = M2V2 for every titration, which is wrong when the balanced equation is not a 1:1 mole ratio.
Practice Questions
- 1 A student titrates 25.00 mL of HCl with 0.1000 M NaOH. The NaOH burette reading changes from 1.20 mL to 26.35 mL. Assuming a 1:1 reaction, what is the molarity of the HCl?
- 2 A 20.00 mL sample of Ca(OH)2 solution is titrated with 0.150 M HCl and requires 18.40 mL of acid. Using Ca(OH)2 + 2HCl -> CaCl2 + 2H2O, calculate the molarity of Ca(OH)2.
- 3 Explain why a back titration may be better than a direct titration for a solid sample that reacts slowly with acid.