Indicators

• Indicators change colour over a narrow pH range approximately centred around the pKa of the indicator.
• Indicators are weak acids
• You only add a few drops of indicator since they are weak acids so they can change the pH value
• Acidic buffer

  HA ⇌ H+ + A-

  When an acid is added,

  • [H+] increases
  • Which should decrease pH but the [H+] reacts with [A-] and the equilibrium shifts to the left
  • Which decreases [H+]
  • So pH stays the same

  When a base is added

  • [OH-] increases
  • It combines with H+ so [H+] decreases
  • Which should increase pH but equilibrium shifts to the right to replace lost [H+]
  • So pH stays almost the same
• Basic buffer

  B + H2O ⇌ BH+ + OH-

  When acid is added

  • [H+] increases which should decrease pH
  • But it combines with OH- decreasing its concentration
  • Equilibrium shifts to the right to replace lost OH-
  • [OH-] increases so pH stays almost the same

  When base added

  • [OH-] increases which should increase pH
  • But the excess OH- reacts with BH+ and the equilibrium shifts left
  • Decreasing [OH-]
  • pH stays almost the same
• Neutralisation point = volume required to neutralise the substance
• The equivalence point would be halfway between the pH curves
• Identify the titration curve

  The equivalence point is around 7

  A) Strong acid

  B) Strong base

  
• Identify titration curve

  Equivalence point is around 9

  A) Weak acid

  B) Strong base

  
• Identify titration curve

  Equivalence point is around 5

  A) Strong acid

  B) Weak base

  
• Identify titration curve

  Equivalence point is around 7

  A) Weak acid

  B) Weak base

  
• In weak acid-weak base titrations, you should not use a pH indicator because there is no sharp, easily detectable change in pH at the equivalence point. This means the indicator's colour change will likely not occur precisely at the equivalence point
• pH at half-neutralisation = pKa so if asked to work out pH when half of the acid has reacted use formula log(Ka)