Question

An electrochemical cell has an Eocell of 1.50 V and transfers 1 mole of electrons. The Faraday constant is 96,485 C/mol e-.
What is the value of delta G?

Answer 1

Answer:

ΔG° = -1.45 × 10⁵ J

Explanation:

Step 1: Given data

• Standard cell potential (E°cell): 1.50 V

• Moles of electrons transferred (n): 1 mol e⁻

• Faraday constant (F): 96,485 C/mol e⁻

Step 2: Calculate the standard Gibbs free energy change (ΔG°)

We will use the following expression.

ΔG° = -n × F × E°cell

ΔG° = -1 mol e⁻ × 96,485 C/mol e⁻ × 1.50 V

ΔG° = -1.45 × 10⁵ J

Answer 2

By apply Gibbs's free energy, the value of delta G is equal to -144727.5 Joules.

Given the following data:

• Faraday constant = 96,485 C/mol e-

• $E^{ \circ}_{cell}$ = 1.50 Volts

• Number of moles of electrons = 1 mol

To determine the value of delta G, we would apply Gibbs's free energy:

Mathematically, Gibbs's free energy is given by the formula:

$\Delta G^\circ = -nFE^{ \circ}_{cell}$

Where:

• F is Faraday constant.

• n is the number of moles.

• $E^{ \circ}_{cell}$ is the electromotive force.

Substituting the given parameters into the formula, we have;

$\Delta G^\circ = -1 \times 96485 \times 1.50$

$\Delta G^\circ = -144727.5$

Delta G = -144727.5 Joules

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