indicators & buffers

• buffers are defined as solutions which resist the change in pH
• buffers remain at a nearly constant pH when adding small amounts of acid or base or by diluting with water
• there are two main types of buffers according to their components

  • acid buffer
  • base buffer
• an acid buffer consists of a mixture of a weak acid and its conjugate base
• in an acid buffer its pH is lower than 7
• a base buffer consists of a weak base and its conjugate acid
• the pH of a base buffer is higher than 7
• a buffer works by setting up an equilibrium in each buffer solution
• the buffering action of an acidic buffer can be represented by

  HA (acid) <--> H+ (conjugate acid) + A- (conjugate base)
• the required composition of an acid buffer solution can be calculated from the desired pH and the acid dissociation constant
• buffers are used to maintain the pH of solutions at a constant value, even when small amounts of acids or bases are added.
• an approximate pH of an acid buffer solution can be calculated by the acid dissociation constant

  pH = pKa - log10 [acid]/[salt]

  or

  pH = pKa + log10 [acid]/[salt]
• in an acid buffer there's lots of weak acid molecules present since it only partly dissociates into the conjugate base and acid
• the addition of a small volume of dilute acid into an acidic buffer will increase the concentration of hydrogen ions in the equilibrium

  the position of the equilibrium is disrupted, shifting it to the left so the acid removes the excess ions
• for a buffer to be able to stabilise the pH of a solution, it must be able to react with the added OH- and H+ ions
• the addition of a small volume of diluted base to the buffer solution increases the concentration of hydroxide ions in the equilibrium

  the OH- ions react with the H+ ions in the buffer, so removing some of them, which causes the equilibrium to move to the right to replace the H+ ions that have reacted with the OH- ions
• an example of an acidic buffer is ethanoic acid and sodium ethanoate
• an example of a basic buffer is ammonia solution and ammonium chloride
• when a small volume of dilute acid is added to a basic buffer, the additional hydrogen ions are removed as they react with the weak base in the solution to form ions

  the equilibrium position moves left
• when small volumes of base is added to a basic buffer, the system tries to counteract the effect by removing the hydroxide ions which react with the hydrogen ions in the buffer to form water

  the equilibrium moves right to replace the removed hydrogen ions
• basic buffer


  • if H+ is added = combines with weak base
  • if OH- is added = combines with conjugate acid (provided by salt)
  • weak base = traps added H+
  • salt = traps OH-
• acidic buffer


  • weak acid = traps OH-
  • salt of acid = traps H+
• indicators are weak acid or bases which are able to give a measure of the pH of a solution by that colour
• when deciding on which indicator to use in a titration the equivalence point in the titration must be considered
• the equivalence point is the point in a titration where enough titrant added is enough to completely neutralise the analyte solution
• the shape of the titration curve indicates the type of titration
• the vertical region in a titration curve indicates where there's a rapid change in pH
• the equivalence point depends on the strength of the acid and base used and the nature of the salt formed at the end-point
• strong acid & strong base


  vertical region occurs between pH 3 and pH 10
• strong acid & weak base


  vertical region occurs between pH 3 and pH 8
• weak acid & strong base


  vertical region occurs between pH 7 and pH 10
• weak acid & weak base


  no sharp vertical region so there's no rapid change in pH, therefore no suitable indicator
• when choosing which indicator to use, a colour-change at a pH equal to the pH of the equivalence point is desired

  this means the colour change occurs in the vertical region of the titration curve
• the general formula for an indicator can be represented by

  HIn (aq) + H+ (aq) <--> H3O+ (aq) + In- (aq)
• In = indicator
• for a good indicator the undissociated acid, HIn, will have a distinctly different colour from its conjugate base In-
• the acid dissociation constant can be determined by the equation

  KIn = [H3O+][In-]/[HIn]

  which can be rearranged to

  [In-][HIN-]=KIn/[H3O-]
• the pH range over which a colour change can be seen is established using

  pH = pKa +/- 1
• the pH of a solution is determined by the pKIn of the indicator and the ratio of [In-] to [HIn]

  since these are different colours, the ratio determines the overall colour of the solution
• for a given indicator, the pH of the solution determines the overall colour