Chem Lab Exam

• Determination of Acid Content in Vinegar


  1. Weigh 3 ml of vinegar and transfer it to a flask

  2. Dilute the vinegar with 50 ml of water in the flask

  3. Add phenolphthalein indicator to the solution

  4. Titrate the solution with the standard NaOH solution

  5. Record the volume of NaOH required to reach the neutralization point
• Vinegar


  • Typically contains a 4-5% acetic acid solution
  • Dilution of the vinegar with water helps in achieving accurate titration results
• Phenolphthalein


  Used as an indicator to detect the endpoint of the titration
• Neutralization point


  Reached when the solution changes color, indicating complete reaction between the acetic acid in vinegar and the NaOH solution
• Acetic acid, the main component of vinegar, is a weak acid that does not dissociate significantly in solution
• The addition of water to the vinegar sample helps in minimizing dilution errors
• Titration with NaOH allows for the determination of the acetic acid content in vinegar based on the volume of NaOH required for neutralization
• This experiment provides a practical method for quantifying the acid content in vinegar, which is essential for various culinary and industrial applications
• Standardization of KMnO4 Solution


  1. Weigh 300 mg of dried reagent grade Na2C2O4 in a tall beaker

  2. Dissolve the sodium oxalate in about 50 mL of 1 M H2SO4

  3. Heat the solution to 80°C and titrate it with KMnO4 while stirring with a thermometer

  4. Determine a blank by titrating an equal volume of 1 M H2SO4

  5. Correct the titration data for the blank

  6. Calculate the concentration of KMnO4 in both molarity and normality
• Sodium oxalate (Na2C2O4)


  • Commonly used as a standard substance in permanganometry due to its well-defined reaction with KMnO4
  • A white crystalline solid with a molecular mass of 134 g/mole
• Reaction between KMnO4 and sodium oxalate

  A redox titration
• KMnO4


  Serves as a self-indicator in acidic solutions
• The reaction proceeds slowly at first but accelerates as manganese(II) accumulates, indicating autocatalysis
• Endpoint


  The expected endpoint is a pale pink color that persists for 30 seconds
• Molarity and normality of KMnO4 are calculated based on the equivalence between KMnO4 and Na2C2O4
• Hydrogen Peroxide Analysis


  1. Weigh a 250 mL Erlenmeyer flask

  2. Pipet 3 mL of agua oxinada into the flask and reweigh

  3. Add 50 mL of 1M H2SO4

  4. Titrate with standard KMnO4 solution until a permanent pale pink color appears

  5. Calculate the concentration of H2O2 in % (w/w), % (w/v), % (v/v), Normality, and Molarity
• Potassium permanganate (KMnO4)


  Oxidizes hydrogen peroxide (H2O2) to water and oxygen
• Reaction between KMnO4 and H2O2


  A redox reaction: 5 H2O2+2MnO4−+6H+→5O2+2Mn2++8H2O
• It's important to use sulfuric acid instead of hydrochloric acid to avoid unwanted side reactions with KMnO4
• Titration Curve


  1. Construct titration curves for various systems

  2. Identify endpoints based on observable physical changes
• Importance of understanding titration curves for accurate analysis
• Titration curves are fundamental tools in analytical chemistry, providing critical insights into acid-base reactions and equilibrium phenomena
• Titration curves


  • Represent the pH changes that occur during acid-base titrations
  • Help determine the equivalence point, endpoint, and buffering capacity of solutions
• Components of a Titration Curve


  • Initial pH
  • Buffer region
  • Equivalence point
  • Beyond the equivalence point
• Factors Affecting Titration Curves


  • Nature of acids and bases
  • Concentration
  • Presence of buffers
• The initial pH of the solution is low (typically around 1) due to the presence of excess H₃O⁺ ions from the strong acid
• Buffer Region


  • As the strong base is added, the pH increases gradually, indicating the formation of a buffer solution
  • The buffer region exhibits a gradual rise in pH due to the presence of both the weak conjugate acid of the strong base and the strong acid
• Equivalence Point


  • The equivalence point is reached when the moles of acid equal the moles of base
  • At this point, the solution is neutral, and the pH is approximately 7
• Beyond Equivalence Point


  • Beyond the equivalence point, the excess strong base leads to a rapid increase in pH
  • The solution becomes basic, and the pH rises sharply due to the excess OH⁻ ions
• The titration curve for the titration of a strong acid with a strong base exhibits the following characteristics: a steep rise in pH in the buffer region, a sharp vertical rise at the equivalence point, a plateau at the equivalence point where the pH remains constant, and a gradual rise in pH beyond the equivalence point due to the excess base