An indicator has the reaction
HIn(aq)<-->+In-(aq). Explain how the indicator can change color when a strong acid is poured into a solution of this indicator.
Did you intend to write
HIn(aq) <==> H^+(aq) + In^-(aq)
addition of H^+ from an acid causes the equilibrium to shift which way? to the left? And this changes the HIn/In^- ratio.
Yes, I apologize for the incorrect representation. The correct balanced equation for the reaction is as follows:
HIn(aq) ⇌ H⁺(aq) + In⁻(aq)
When a strong acid is poured into a solution of this indicator, it introduces an excess of H⁺ ions into the solution. According to Le Chatelier's principle, an increase in the concentration of one of the reactants will shift the equilibrium in the direction that consumes that reactant.
In this case, the addition of H⁺ ions from the strong acid causes the equilibrium to shift to the left. This means that more HIn molecules will react with the additional H⁺ ions to form more undissociated HIn. Consequently, the concentration of In⁻ ions will decrease as they are consumed in the reverse reaction.
The color change of the indicator is due to the fact that the protonated form, HIn, and the deprotonated form, In⁻, have different colors. Typically, HIn is a different color from In⁻. So, when the equilibrium shifts to the left, there is a higher concentration of HIn compared to In⁻, resulting in a color change towards the color of HIn.
To summarize, the addition of a strong acid increases the concentration of H⁺ ions, causing the equilibrium to shift to the left, resulting in a higher concentration of HIn and a color change towards the color of HIn.