Acid-base titrations are a crucial analytical method in chemistry. By gradually adding a known solution to an unknown one, we can determine concentrations and reach the equivalence point where neutralization occurs.

Indicators play a key role in titrations, changing color at specific pH ranges. Understanding how to choose the right indicator and interpret color changes is essential for accurate results in various applications, from food production to environmental monitoring.

Acid-Base Titrations

Process of acid-base titration

• Quantitative analytical method determines concentration of unknown acid or base solution
• Gradual addition of standard solution (titrant) of known concentration to unknown solution (analyte)
• Titrant and analyte react in neutralization reaction
• Equivalence point reached when analyte completely neutralized by titrant
• Applications include determining concentrations in various solutions (vinegar, fruit juices, cleaning products), quality control in industrial processes (food and beverage production, water treatment), and environmental monitoring (measuring acidity of rain or soil samples)

Indicators for acid-base titrations

• Weak acids or bases change color at specific pH ranges
• Color change occurs when indicator's structure shifts between protonated and deprotonated forms
• Choice of indicator depends on expected pH at equivalence point of titration
  • Indicator should have pKa value close to pH at equivalence point
  • Color change should be distinct and easily observable
• Common indicators and their pH ranges:
  • Methyl orange: pH 3.1-4.4 (red to yellow)
  • Bromocresol green: pH 3.8-5.4 (yellow to blue)
  • Phenolphthalein: pH 8.2-10.0 (colorless to pink)

Calculations and Interpretation

Concentration calculations from titration data

• Calculate concentration of unknown solution using equation: $M_a \times V_a = M_b \times V_b$
  • $M_a$: Molarity of acid
  • $V_a$: Volume of acid
  • $M_b$: Molarity of base
  • $V_b$: Volume of base
• Steps to calculate concentration:
  1. Record initial and final volumes of titrant (standard solution) used
  2. Calculate volume of titrant consumed by subtracting initial volume from final volume
  3. Substitute known values into equation and solve for unknown concentration

Color changes vs pH in titrations

• pH of solution changes gradually as titrant added to analyte
  • Indicator's color changes at specific pH ranges, indicating progress of titration
• Before equivalence point:
  • In acid-base titration, pH changes slowly as titrant added
  • Indicator's color will be that of its protonated form
• At equivalence point:
  • pH changes rapidly, causing sharp color change in indicator
  • Color will be intermediate shade between protonated and deprotonated forms
• After equivalence point:
  • pH continues to change slowly as excess titrant added
  • Indicator's color will be that of its deprotonated form
• Observing color changes and relating them to pH helps determine when equivalence point reached and titration complete