The rapid change of pH near the stoichiometric point of an acid-base titration is the basis of indicator detection. pH of the solution is related to ratio of the concentrations of the conjugate acid (HIn) and base ($In^-$) forms of the indicator by the expression -

An acid-base indicator is a weak acid: $HIn \rightleftharpoons H^+ + In^-$
$K_{In} = \dfrac{[H^+][In^-]}{[HIn]}$
Taking negative logarithms: $-\log K_{In} = -\log[H^+] - \log\dfrac{[In^-]}{[HIn]}$
$pK_{In} = pH - \log\dfrac{[In^-]}{[HIn]}$
$\Rightarrow \log\dfrac{[In^-]}{[HIn]} = pH - pK_{In}$ (Henderson equation for the indicator)
Multiplying both sides by $-1$ and using $\log\dfrac{[HIn]}{[In^-]} = -\log\dfrac{[In^-]}{[HIn]}$:
$\log\dfrac{[HIn]}{[In^-]} = pK_{In} - pH$
Both forms are the same relation written two ways, so both $\log\dfrac{[In^-]}{[HIn]} = pH - pK_{In}$ and $\log\dfrac{[HIn]}{[In^-]} = pK_{In} - pH$ are correct.
Note: the answer key of the source accepts both of these expressions as correct, so both are marked correct here.