(481) Unit 4.6 - Introduction to Titration

Chapter 1: Introduction

• Welcome to Unit 4.6

• Introduction to the concept of titration.

Chapter 2: Titration

• Purpose of Titration: To determine the concentration of an unknown substance.

• Components:

  • Titrant: Known concentration solution, located in the burette.

  • Analyte: Unknown substance in the flask (the substance being analyzed).

• Equivalence Point:

  • Also known as the stoichiometric point.

  • Achieved when enough titrant reacts with the analyte.

  • At this point, the moles of titrant equal the moles of analyte based on stoichiometry.

  • Often indicated by a color change due to an indicator in acid-base titrations.

• Key Terminology:

  • Endpoint: The point at which the indicator changes color, synonymous with the equivalence point.

• Steps in a Titration Process:

  1. Unknown Acid in the flask.

  2. Add a few drops of indicator.

  3. Record initial volume of titrant in the burette.

  4. Begin adding titrant to the analyte.

  5. Observe color change indicating equivalence point has been reached.

  6. Calculate the volume of titrant used by subtracting the initial reading from the final.

Chapter 3: Example

• Example of Acid-Base Titration:

  • Hydrofluoric acid reacts with sodium hydroxide.

  • Molecular Equation: HF + NaOH → NaF + H2O

  • Identify which substances ionize based on solubility rules:

    • HF: Weak acid (does not ionize).

    • NaOH: Strong base (ionizes completely).

    • NaF: Ionizes.

    • H2O: Remains intact.

  • Net Ionic Equation:

    • HF + OH⁻ → F⁻ + H2O

• Monitoring pH During Titration:

  • pH curve shows a characteristic S-shape.

  • Find the equivalence point at the most vertical point on the curve.

• Stoichiometry in Titration Calculations:

  • Given: 40 mL of 2 M HF, asked to find mL of 3 M NaOH needed at the equivalence point.

  • Convert 40 mL to liters: 0.04 L.

  • Calculate moles of HF: molarity x volume.

  • Use stoichiometry to relate moles of HF to OH⁻, then to NaOH.

• Next Example:

  • 10 mL of 2 M acetic acid titrated with NaOH (unknown concentration).

  • Write net ionic equation and use the titration curve for calculations (equivalence point at ~14 mL).

  • Molar concentration of NaOH calculated using moles of acetic acid at equivalence point.

Chapter 4: Redox Titrations

• Not all titrations are acid-base; they can be redox reactions.

• Redox Titration: Involves oxidation-reduction reactions; indicators are not always necessary.

  • Example: Oxalic acid titrated with potassium permanganate (dark purple color).

• Balanced Equation helps in stoichiometric calculations:

  • Calculate moles of permanganate and determine molarity of oxalic acid.

  • The coefficients in redox reactions may differ from those in acid-base but processes remain consistent.

• Learning Outcome: Introduction to titration principles, stoichiometry application in both acid-base and redox titrations.

• Further calculations on acid-base titrations will be covered later in the unit.