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I am reading about indicators and their behaviour as weak acids.

I know that the end-point is when exactly all of the H+ ions from an acid have neutralised the OH- ions of the base.

The book later says

At the end point of a titration, the indicator contains equal concentrations of HA and A- and the colour will be inbetween the two extreme colours e.g. for methyl orange, the colour at its end point is orange.

It continues to say

For an indicator HA ⇌ $H^+$ + $A^-$
At the end-point, [HA] = [$A^-$] and $$K_a = \frac{[H^+][A^-]}{[HA]} = [H^+]$$

I am confused why this is true. why is the expression [HA] = [$A^-$] valid?

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In an Acid titration, we want the end-point to be as close to the equivalence point as possible.

Different indicators have different end-points (they will change colour at different pH values).

When an end-point is reached, the colour of the indicator will change with respect to the [$H^+$] concentration of the solution it is measuring.

As [$H^+$] increases, there will be a point when [HA] = [$A^-$] and depending on the indicator used by the person doing the titration, this would be close to the calculated equivalence-point.

Note, the difference in volume between the end-point and equivalence-point is very very small, hence we can approximate the end-point as the same as the equivalence-point in any experiment.

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