Determining the nature of common pH indicators (methyl orange) using Le Chatlier's principle I got some good information about this topic yesterday, and I understood what I was told, but I have come up with a specific way of phrasing the problem I'm having. It is slightly different than before, and I'm looking for someone to tell me if what I'm doing makes any sense or not, rather than just rote learning facts like 'methyl orange indicator is a weak acid/base'. I want to be able to use this process for any weakly acidic/basic indicator I come across. 

(Note: Methyl orange is taken to be orange when it has a moderate pH (HIn, not very acidic or basic, and red between a pH of 4.4 and 3.1 - i.e. when it is In- or In+ depending on the scenario below)

Assume methyl orange, an indicator, is a weak acid. It dissociates as follows in aqueous solution;



We know that if an acid is added to this system at equilibrium, the concentration of hydrogen ions will increase. Using Le Chatlier's principle; to oppose this stress, the reaction shifts to the right, and the colour of the initial indicator (in this case methyl orange, so orange) will dominate and become more intense.

The solution will not go red, like it would if a base was added and proton concentration was decreased, as the equilibrium would shift in the other direction.

The problem is, that if you add an acid to this system in real life, a colour change occurs at a pH of 4.4, from orange to red. This means that methyl orange indicator cannot be an acid. Right?

If you instead assume that methyl orange is a weak base, it dissociates like so in aqueous solution; 



So here, if you add an acid, the concentration of hydroxide ions will decrease (due to bonding between the hydrogen ions in the acid and these hydroxide ions to form water), so the equilibrium shifts to the right, to compensate.

This means that the red colour will dominate.

This result is in line with observation, as methyl indicator will change colour from orange to red at a pH of 3.1-4.4, which is acidic.

Does my reasoning therefore demonstrate that methyl indicator must be a weak base?
Methyl Orange is the common name for Sodium 4-[(4-dimethylamino)phenyldiazenyl]benzenesulfonate. So the compound is a salt or a conjugate base. It is the sodium salt for the indicator since that form is more soluble than the acid form. 

 

The pKa is 3.47 in aqueous solution, so at pH values of 3.1 or lower the molecule is protonated.



At pH values greater than 4.4 all the molecules in solution are deprotonated as the sulfonate anion. 



Linked two images from ChemWiki.



The confusion is caused by rather poor wording in a question.


I'd prefer the wording to be:


  
    
      ... In aqueous solution methyl orange indicator behaves like a weak acid. Letting HX represent the protonated form of the indicator, it dissociates as follows:
    
  


It is splitting hairs, but methyl orange is the name of the sodium salt, not the protonated form of the indicator.