In the given electrochemical cell, Ag(s)|AgCl(s)| FeCl_ 2 (aq),...

The cell Ag(s)|AgCl(s)|FeCl₂(aq), FeCl₃(aq)|Pt(s) has the reaction:
Ag(s) + Cl⁻ + Fe³⁺ → AgCl(s) + Fe²⁺
Using Nernst equation:
$E_{cell} = E°_{cell} - \frac{0.059}{n}\log Q$ where $Q = \frac{[Fe^{2+}]}{[Fe^{3+}][Cl^-]}$
For $E_{cell}$ to increase, Q must decrease. Analyzing each condition:
A. Increase [Fe²⁺] increases Q, decreases $E_{cell}$ - false
B. Decrease [Fe³⁺] increases Q (smaller denominator), decreases $E_{cell}$ - false
C. Decrease [Fe²⁺] decreases Q, increases $E_{cell}$ - true
D. Increase [Fe³⁺] decreases Q, increases $E_{cell}$ - true
E. Increase [Cl⁻] decreases Q, increases $E_{cell}$ - true
Therefore C, D, and E cause $E_{cell}$ to increase.