The cell notation below represents an electrochemical cell: Cu/Cu^{2+} (1 mol·dm^{-3}) // Ag^{+} (1 mol·dm^{-3}) / Ag 298 K/25 °C 6.1.1 What energy conversion is taking place in the above cell? 6.1.2 Write down TWO indicators from the cell notation that prove that the cell is operating under standard conditions - NSC Technical Sciences - Question 6 - 2022 - Paper 2

1. The concentration of the electrolytes is specified as 1 mol·dm^{-3}, indicating standard concentration.
2. The temperature is given as 298 K (25 °C), which is standard temperature.

The balanced half-reaction that occurs at the cathode is:

Ag^{+}(aq) + e^{-} → Ag(s)

The balanced half-reaction that occurs at the anode is:

Cu(s) → Cu^{2+}(aq) + 2e^{-}

To determine whether the reaction is spontaneous, calculate the cell potential ( E_{cell}^{ heta}) using:

$E_{cell}^{ heta} = E_{cathode}^{ heta} - E_{anode}^{ heta}$

Given that:

• $E_{cathode}^{ heta} = +0.80 V$ (for Ag)
• $E_{anode}^{ heta} = +0.34 V$ (for Cu)

We have:

$E_{cell}^{ heta} = 0.80 V - 0.34 V = 0.46 V$

Since the cell potential is positive, the reaction is spontaneous.

The reason the reaction is spontaneous is that the positive cell potential ( E_{cell}^{ heta} = 0.46 V) implies that the Gibbs free energy change ( ΔG) for the reaction is negative, following the relationship $ΔG = -nFE_{cell}^{ heta}$. A negative Gibbs free energy indicates a spontaneous reaction.