Chapter 17, Problem 92QAP

Interpretation Introduction

Interpretation:

The voltage produced when CO reacts with O₂ to produce CO₂ needs to be calculated given that all the gases are at a pressure of 1 atm and at a temperature of 1000 °C.

Concept introduction:

The change in the Gibbs free energy is that Δ G is a thermodynamic function which governs the spontaneity of a chemical reaction. If Δ G is negative the reaction is spontaneous, positive value indicated that the reaction is non-spontaneous and if Δ G = 0, then the reaction is said to be at equilibrium.

The standard Gibbs free energy ${\text{ΔG}}^{\text{0}}$ for a given chemical reaction can be expressed as a function of temperature, T via the Gibbs-Helmholtz equation:

${\text{ΔG}}^{\text{0}}{\text{ = ΔH}}^{\text{0}}{\text{ - TΔS}}^{\text{0}}\text{ -------(1)}$

where, ${\text{ΔH}}^{\text{0}}$ is the standard enthalpy change, and ${\text{ΔS}}^{\text{0}}$ is the standard entropy change

It is also related to the standard voltage (E°) by the equation:

$\begin{array}{l}{\text{ΔG}}^{\text{0}}{\text{ = -nFE}}^0\text{ -------(2)}\\ \text{where n = number of electrons, F = Faraday constant (96500 C)}\end{array}$

The cell voltage under non-standard conditions (E) is related to the standard voltage (E°) via the Nernst equation:

$\begin{array}{l}{\text{E = E}}^{\text{0}}\text{ - }\frac{\text{0}\text{.0257}}{\text{n}}\text{lnQ --------(3)}\\ \text{where n = number electrons involved in the redox reaction}\\ \text{Q = reaction quotient = }\frac{\text{[Products]}}{\text{[Reactants]}}\end{array}$