A galvanic cell is set up at 298 K as Ag(s) | AgCl(s), KCl(0.1 M) \| AgNO3(0.1 M) | Ag(s). Its measured EMF is 0.472 V. Take 2.303RT/F=0.059 V and…

Both electrodes are silver, so this is a concentration cell with the cathode (right) having $[\mathrm{Ag^+}]=0.1\ \mathrm{M}$ and the anode (left) having $[\mathrm{Ag^+}]$ fixed by the solubility equilibrium of AgCl in $0.1\ \mathrm{M}$ chloride: $[\mathrm{Ag^+}]_L=K_{sp}/0.1$. Nernst equation: $E=0.059\log\dfrac{[\mathrm{Ag^+}]_R}{[\mathrm{Ag^+}]_L}=0.059\log\dfrac{0.1\times0.1}{K_{sp}}$. Setting $E=0.472$ gives $\log(0.01/K_{sp})=8$, so $K_{sp}=10^{-10}$. In pure water $s=\sqrt{K_{sp}}=1.0\times10^{-5}\ \mathrm{M}$. Using $[\mathrm{Ag^+}]_L=K_{sp}$ (ignoring common ion) gives $3.2\times10^{-5}$; putting the chloride concentration on the wrong side gives $10^{-4}$; quoting $K_{sp}$ itself gives $10^{-10}$. Hence the answer is $1.0\times10^{-5}\ \mathrm{M}$. As a check, the half-cell on the left behaves as a chloride electrode: its potential is lower than that of the silver electrode by exactly $0.472$ V, consistent with $[\mathrm{Ag^+}]_L=10^{-9}\ \mathrm{M}$.