Summary
Highlights
The lecture introduces module 6 on acid-base or neutralization titration, a volumetric analysis method. Unlike gravimetric analysis which focuses on mass, volumetric analysis measures analyte content by the volume of titrant added. Four types of titration are mentioned: neutralization, complexometric, redox, and precipitation. This lecture exclusively covers neutralization titration, based on acid-base reactions, using an outline including volumetric methods, titration setup, neutralization concepts, titrants, primary standards, acid-base indicators, and calculation examples.
Volumetric analysis involves titration, adding a titrant from a buret to a sample to determine the analyte's amount based on the volume required to reach an equivalence point. The typical setup involves the titrant in a buret and the analyte in an Erlenmeyer flask, with proper handling demonstrated. Neutralization titration is broadly used in chemistry to determine inorganic and organic compounds with sufficient acidity or basicity, using strong acids and strong bases as titrants. The sample can be a strong or weak acid/base, but the titrant must always be a strong acid or base.
Two types of titrants are discussed: base titrants (most commonly sodium hydroxide, which is a secondary standard due to hygroscopy and impurities, requiring standardization) and acid titrants (typically HCl, also a secondary standard). Primary standards are ultra-pure compounds used as reference materials for titrimetric analysis, requiring high purity, stability, solubility, low cost, absence of hydrate water, and a large molar mass. Common primary standards for base titrants include potassium acid phthalate (KHP), 2,4,6-trinitrobenzoic acid, and benzoic acid. For acid titrants, sodium carbonate and sodium tetraborate are common.
Titration basics involve the titrant, analyte, and indicators. Back titration is a technique used when the analyte reacts slowly, involving an excess standard agent followed by titration with a secondary standard. Acid-base indicators are weak organic acids or bases that change color based on pH, acting as secondary acids or bases and being weaker than the analyte. Examples include Bromophenol Blue, Congo Red, Methyl Orange, and Phenolphthalein. The indicator's pH range, where color change occurs, is typically pKa ± 1.
Titration curves plot pH against the volume of titrant added, typically showing a sigmoidal shape. The most suitable indicator has a pKa closest to the pH at the equivalence point. Errors in acid-base titration include determinate errors (indicator color change pH differs from equivalence point pH) and indeterminate errors (e.g., color blindness affecting observation). The curve has four stages: before titrant addition, pre-equivalence point, equivalence point (where stoichiometry is reached and indicator changes color), and post-equivalence point (excess titrant).
An example calculation is provided for titrating 50 mL of 0.1 M HCl with 0.1 M NaOH. The pH at various points (start, after 10 mL NaOH, at equivalence point, and after 60 mL NaOH) is calculated. Before addition, pH is 1 (from HCl dissociation). After 10 mL NaOH, the solution is still acidic, and pH increases to 1.18. At the equivalence point (50 mL NaOH), the pH is 7, as only NaCl and water are present. After 60 mL NaOH (post-equivalence point), the solution becomes basic with a pH of 11.6.
The calculation for titrating a weak acid (50 mL of 0.1 M acetic acid) with a strong base (0.1 M NaOH) is explained. Before titration, the pH is calculated using the acid dissociation constant (Ka) of acetic acid, resulting in a pH of 2.87. At the pre-equivalence point after adding 10 mL of NaOH, the Henderson-Hasselbalch equation is used with concentrations of acetic acid and acetate ion, yielding a pH of 4.14. At halfway to the equivalence point (25 mL NaOH), the pH equals the pKa, which is 4.74. At the equivalence point (50 mL NaOH), the solution is treated as a weak base, and the pH is 8.87. In the post-equivalence point, the pH calculation follows the strong base logic as in the previous example, resulting in a pH of 11.6.
The discussion briefly mentions the titration of weak bases with strong acids, where the curve would descend. It concludes by stating that titrating a weak acid with a weak base is not feasible for obtaining a clear endpoint.