Explanation

The overall cell reaction is governed by the half reactions:

At Ag/AgCl electrode:

$$ \text{AgCl(s)} + e^- \rightleftharpoons \text{Ag(s)} + \text{Cl}^- $$

At platinum electrode:

$$ \text{Fe}^{3+} + e^- \rightleftharpoons \text{Fe}^{2+} $$

The Nernst equation for the cell potential can be written as

$$ E_{\text{cell}} = E^\circ_{\text{cell}} - \frac{0.059}{1}\log\!\left(\frac{[\text{Fe}^{2+}][\text{Cl}^-]}{[\text{Fe}^{3+}]}\right) $$

From the expression, $E_{\text{cell}}$ increases when the reaction quotient decreases.

Now check each condition:

Increasing Fe$^{2+}$ increases the reaction quotient, hence $E_{\text{cell}}$ decreases (A is false).

Decreasing Fe$^{3+}$ increases the reaction quotient, hence $E_{\text{cell}}$ decreases (B is false).

Decreasing Fe$^{2+}$ decreases the reaction quotient, hence $E_{\text{cell}}$ increases (C is true).

Increasing Fe$^{3+}$ decreases the reaction quotient, hence $E_{\text{cell}}$ increases (D is true).

Increasing Cl$^-$ decreases the electrode potential of Ag/AgCl electrode in such a way that overall cell potential increases (E is true).

Therefore, the correct conditions are

$$ \boxed{\text{C, D and E only}} $$

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