Most indicators are weak acids where the indicator, HIn, is one colour and its conjugate base, In^–, is a different colour. The following video shows what happens when the indicator, bromothymol blue, is subjected to acidic or alkaline conditions.

INSERT [IND_CI_01_vid  or  IND_CI_01_ima if not video, or could be a simple animation with voice commentary explaining what's happening and the equilibrium drawn out]

[video requested here if sufficient budget left as this shows the equilibrium and draws on the points made in the last 3 topics about equilibria. However, the image shows the simple colour change but not the reversible equilibrium aspects directly].

The equilibrium for an indicator, HIn, is represented by the equilibria:

HIn ∏  H⁺ + In^-  

Or equally, represted as:  HIn_(aq) + H₂O_(l)  ∏  H₃O⁺_(aq) + In^- _(aq)            

When acid is added to a solution of bromothymol blue, where does the position of equilibrium lie and what is the main species present in solution?

HIn (yellow)  ∏  H⁺ + In^-  (blue)

In acidic solution, the equilibrium shifts to the left and HIn predominates. This is yellow. Adding acid increases the concentration of H⁺ in solution and the ‘position of equilibrium’ shifts in the direction that reduces the concentration of hydrogen ions, ie to the left, and HIn predominates. The indicator turns  yellow.

When bromothymol blue is in alkaline solution, where does the position of equilibrium lie and what is the main species present in solution?

In alkaline solution, the equilibrium shifts to the right and In^– predominates. This is blue.

The dissociation constant of the acid indicator, KIn, is expressed as: K_In = [H₃O⁺][In^–] / [HIn] What can you say about the pH of the colour change?

During the colour change, halfway, when [H₃O⁺] = [In^–] these concentrations cancel from the equation to leave K_In = [H₃O⁺]. So in theory, the colour change occurs when K_In = [H₃O⁺] or, taking logs, pK_In = pH