Question

Consider the following reaction at constant pressure. Use the information provided below to determine the value of ΔS at 473 K. Predict whether or not this reaction will be spontaneous at this temperature.
4NH3 (g) + 3O2 (g) → 2N2 (g) + 6H2O (g) ΔH = –1267 kJ

Answer 1

Answer:

The reaction will be spontaneous

Explanation:

To determine if the reaction will be spontaneous or not at this temperature, we need to calculate the Gibbs's energy using the following formula:

$\Delta G= \Delta H - T * \Delta S$

If the Gibbs's energy is negative, the reaction will be spontaneous, but if it's positive it will not.

Calculating the $\Delta G= -1267 - 473 K* \Delta S$ :

$\Delta G= -1267 - 473 K* \Delta S$

Now, other factor we need to determine is the sign of the S variation. When talking about gases, the more moles you have in your system the more enthropic it is.

In this reaction you go from 7 moles to 8 moles of gas, so you can say that you are going from one enthropy to another higher than the first one. This results in: $\Delta S>0[/tex}

Back to this expression:

[tex]\Delta G= -1267 - 473 K* \Delta S$

If the variation of S is positive, the Gibbs's energy will be negative always and the reaction will be spontaneous.