   Chapter 21, Problem 62PS

Chapter
Section
Textbook Problem

Use data in Appendix L to calculate the enthalpy and free energy change for the reaction 2 NO 2 ( g ) → N 2 O 4 (g) Is this reaction exothermic or endothermic? Is the reaction product- or reactant-favored at equilibrium?

Interpretation Introduction

Interpretation:

To calculate the enthalpy and free energy change for the given reaction and check whether the reaction is exothermic or endothermic.

Concept introduction:

The Gibbs free energy or the free energy change is a thermodynamic quantity represented by ΔG. The expression for the free energy change is:

ΔrG°=nΔfG°(products)nΔfG°(reactants)

The enthalpy change is expressed by the formula,

ΔrH°=nΔfH°(products)nΔfH°(reactants)

If the value of enthalpy change is negative, the reaction is exothermic.

If the value of free energy change is negative, then the reaction is product-favored at equilibrium.

Explanation

The standard free energy change and enthalpy change for the formation of N2O4(g) is calculated below.

Given:

Refer to Appendix L for the values of standard enthalpy and free energy.

The standard enthalpy change value of N2O4(g) is 9.08 kJ/mol.

The standard enthalpy change value of NO2(g) is 33.1 kJ/mol.

The standard free energy change value of N2O4(g) is 97.73 kJ/mol.

The standard free energy change value of NO2(g) is 51.23 kJ/mol.

The given reaction is,

2NO2(g)N2O4(g)

The expression for enthalpy change is,

ΔrH°=nΔfH°(products)nΔfH°(reactants)=[(1 mol N2O4(g)/mol-rxn)ΔfH°[N2O4(g)](2 mol NO2(g)/mol-rxn)ΔfH°[NO2(g)] ]

Substitute the values,

ΔrH°=[(1 mol N2O4(g)/mol-rxn)(9

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