The rapid change of `pH` near the stoichiometric point of an acid-base titration is the basic of indicator detection. `pH` of the solution is related to the ratio of the concentration of conjugate acid `(Hin)` and base `(In^(-))` forms of the indicator by the expression

Options

A. `log.([In^(-)])/([HIn])=pK_(In)-pH`
B. `log.([HIn])/([In^(-)])=pK_(In)-pH`
C. `log.([HIn])/([In^(-)])=pH-pK_(In)`
D. `log.([In^(-)])/([HIn])=pH-pK_(In)`

Answer

D

Explanation

Acid indicatiors are generally weak acid. The dissociation of indicator Hin takes place as floows
`HIn hArr H^(+)+In^(-)`
`therefore K_(In)=([H^(+)][In^(-)])/([HIn])`
or `[H^(+)]=K_(In).([HIn])/([In^(-)])` ….(i)
`because pH=-log [H^(+)]` ...(ii)
From eq. (i) and (ii) we get,
`therefore pH=-log(K_(In).([HIn])/([In^(-)]))`
`=-log K_(In)+log.([In^(-)])/([HIn])=pK_(In)+log.([In^(-)])/([HIn])`
or `log.([In^(-)])/([HIn])=pH-pK_(In)`