The potential in an electrochemical cell, E, is related to the Gibb's free energy change (ΔG) for the overall cell redox reaction: (1) ΔG0 = - n F E0 where n is the number of electrons transferred during the redox reaction, F is Faraday's constant (96,500 C / mol), and the superscript 0 indicates standard conditions (1 atm, 1 M concentrations, and 25 °C). Thus, a measurement of the cell voltage at standard conditions can be used to determine ΔG0. As an example, the following cell reaction: Zn (s) + Cu2 + (aq) → Zn2 + (aq) + Cu (m) generates a cell voltage of + 1.10 V under standard conditions. Calculate and enter delta G degree (with 3 sig figs) for this reaction in kJ/mol.

Question

Jillian Grimes

Chemistry

29 March, 23:36

The potential in an electrochemical cell, E, is related to the Gibb's free energy change (ΔG) for the overall cell redox reaction: (1) ΔG0 = - n F E0 where n is the number of electrons transferred during the redox reaction, F is Faraday's constant (96,500 C / mol), and the superscript 0 indicates standard conditions (1 atm, 1 M concentrations, and 25 °C). Thus, a measurement of the cell voltage at standard conditions can be used to determine ΔG0. As an example, the following cell reaction: Zn (s) + Cu2 + (aq) → Zn2 + (aq) + Cu (m) generates a cell voltage of + 1.10 V under standard conditions. Calculate and enter delta G degree (with 3 sig figs) for this reaction in kJ/mol.

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Parker Liu · Answer 1

As an example, the following cell reaction: Zn (s) + Cu2 + (aq) → Zn2 + (aq) + Cu (m) generates a cell voltage of + 1.10 V under standard conditions. Calculate and enter delta G degree (with 3 sig figs) for this reaction in kJ/mol.

Zn (s) + Cu2 + (aq) → Zn2 + (aq) + Cu (m)

ΔG = ΔG° + RTInQ

Q = 1

ΔG = ΔG°

ΔG = = nFE°

n=no of electrons transfered.

E° = 1.1v

ΔG° = - 2 * 96500 * 1.10

= - 212300J

ΔG° = -212.3kJ/mol

Therefore, the ΔG° = - 212.3kJ/mol