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.
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.