E naught cell for the following galvanic cell is +0.254V.
Hg^2+(aq) + 2I^-(aq) ==> 2 Hg(l) +I2(s)
What is the change G naught for this reaction?
To calculate the change in Gibbs free energy (ΔG°) for the reaction, we can use the formula:
ΔG° = -nFΔE°
Where:
ΔG° is the change in Gibbs free energy
n is the number of moles of electrons transferred in the reaction
F is Faraday's constant (96,485 C/mol)
ΔE° is the standard cell potential (also known as E° cell or E° cell reaction)
From the given reaction: Hg^2+(aq) + 2I^-(aq) ==> 2 Hg(l) + I2(s)
We can see that 2 moles of electrons are transferred in the balanced equation.
Given:
ΔE° = +0.254 V
Thus, substituting the values into the formula:
ΔG° = -(2 mol)(96,485 C/mol)(+0.254 V)
By multiplying the values, we get:
ΔG° = -49,122 J
Therefore, the change in standard Gibbs free energy (ΔG°) for this reaction is -49,122 J or -49.122 kJ.
To determine the standard change in Gibbs free energy (ΔG°) for a reaction, we need to use the Nernst equation, which relates the standard cell potential (E°cell) of the reaction to ΔG°. The Nernst equation is as follows:
ΔG° = -nF E°cell
Where:
ΔG° = standard change in Gibbs free energy
n = number of moles of electrons transferred during the reaction
F = Faraday's constant (which is approximately equal to 96,485 C/mol)
E°cell = standard cell potential
In this case, the balanced equation indicates that 2 moles of electrons are transferred in the reaction. Therefore, n = 2.
Given that E°cell = +0.254 V and n = 2, we can now calculate the value of ΔG°:
ΔG° = -2 * 96485 * 0.254
Solving this equation will give us the value of ΔG°.