Question
The major industrial source of hydrogen gas is by the following reaction:$$\mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \longrightarrow \mathrm{CO}(g)+3 \mathrm{H}_{2}(g)$$Use bond energies to predict $\Delta H$ for this reaction.
Step 1
The bond energies are as follows: C-H: 414 kJ/mol O-H: 463 kJ/mol C-O: 358 kJ/mol H-H: 436 kJ/mol Now, let's calculate the total bond energy for the reactants: 1 C-H bond in CH4: 1 * 414 kJ/mol = 414 kJ/mol 2 O-H bonds in H2O: 2 * 463 kJ/mol = 926 kJ/mol Total Show more…
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The major industrial source of hydrogen gas is by the following reaction: $$ \mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \longrightarrow \mathrm{CO}(g)+3 \mathrm{H}_{2}(g) $$ Use bond energies to predict r this reaction.
Hydrogen is prepared from natural gas (mainly methane, $\mathrm{CH}_{4}$ ) by partial oxidation. $$ \mathrm{CH}_{4}(g)+\frac{1}{2} \mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}(g)+2 \mathrm{H}_{2}(g) $$ Calculate the enthalpy change $\Delta H^{\circ}$ for this reaction, using standard enthalpies of formation.
Hydrogen is prepared from natural gas (mainly methane, $\mathrm{CH}_{4}$ ) by partial oxidation. $$ 2 \mathrm{CH}_{4}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}(g)+4 \mathrm{H}_{2}(g) $$ Calculate the enthalpy change $\Delta H^{\circ}$ for this reaction, using standard enthalpies of formation.
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