Question

The chemical reaction that causes iron to corrode in air is given by 4fe+3o2→2fe2o3 in which at 298 k δh∘rxn = −1684 kj δs∘rxn = −543.7 j/k

Answer 1

Answer:

A. $\Delta G=-1486.0kJ$

B. $\Delta G=301.6kJ$

C. $T=3031.8K$

Explanation:

Hello,

In this case, the requirements, followed by the solutions are:

A . What is the standard Gibbs free energy for this reaction? Assume the commonly used standard reference temperature of 298 K

In this case, given the enthalpy and entropy change for iron's corrosion, we compute the Gibbs free energy as shown below:

$\Delta G=\Delta H- T\Delta S\\\Delta G=-1684kJ -298K*(-543.7\frac{J}{K}*\frac{1kJ}{1000J})\\ \Delta G=-1486.0kJ$

B. What is the Gibbs free energy for this reaction at 3652K ? Assume that ΔH and ΔS do not change with temperature

$\Delta G=\Delta H- T\Delta S\\\Delta G=-1684kJ -3652K*(-543.7\frac{J}{K}*\frac{1kJ}{1000J})\\ \Delta G=301.6kJ$

C. At what temperature Teq do the forward and reverse corrosion reactions occur in equilibrium?

In this case, at equilibrium, the Gibbs free energy of the reaction is zero, therefore, the temperature will be:

$\Delta G=\Delta H- T\Delta S\\0=-1684kJ -T*(-543.7\frac{J}{K}*\frac{1kJ}{1000J})\\T=\frac{-1684kJ}{-0.5437\frac{kJ}{K}} \\\\T=3031.8K$

Best regards.