A student investigates the relationship between ion concentration and cell potential by constructing a series of \(\text{Ag/Ag}^+\) concentration cells at \(298 \text{ K}\). Each cell consists of two \(\text{Ag(s)}\) electrodes in separate beakers connected by a salt bridge and a voltmeter. The cathode compartment contains \([\text{Ag}^+] = 1.0 \text{ M}\) in every trial, while the anode compartment contains \(\text{AgNO}_3\text{(aq)}\) at various concentrations. The initial cell potential, \(E_{\text{cell}}\), recorded for each trial is shown in the table below.

Trial	\([\text{Ag}^+]_{\text{anode}}\) (M)	\([\text{Ag}^+]_{\text{cathode}}\) (M)	\(E_{\text{cell}}\) (V)
1	\(1.0 \times 10^{-1}\)	\(1.0\)	\(+0.059\)
2	\(1.0 \times 10^{-2}\)	\(1.0\)	\(+0.118\)
3	\(1.0 \times 10^{-3}\)	\(1.0\)	\(+0.177\)
4	Unknown \(X\)	\(1.0\)	\(+0.236\)

Based on the data, what is \([\text{Ag}^+]\) in the unknown solution, and what condition explains why \(E_{\text{cell}}\) becomes \(0.000 \text{ V}\) after the cell is allowed to operate for a long period of time?