Examine the given reaction. NH4Cl(s) → NH3(g) + HCl(g)
∆H0=176 kJ/mol
∆S0=0.285 kJ/(mol K)
What can be deduced from to this reaction?

A) At room temperature the reaction is not spontaneous. However, at lower temperatures, like –40 °C, the free energy value turns negative and this reaction becomes spontaneous.
B) At room temperature, the reaction is spontaneous. However, at high temperatures, like 800 °C, the free energy value turns negative and this reaction becomes nonspontaneous.
C) At room temperature, the reaction is spontaneous. If the temperature is increased, the free energy value turns negative and the reaction stays spontaneous.
D) At room temperature, the reaction is not spontaneous. However, at high temperatures, like 800 °C, the free energy value turns negative and this reaction becomes spontaneous.

The correct option is;

D)

Explanation:

The given reaction is presented as follows;

NH₄Cl (s) → NH₃ (g) + HCl (g) ΔH° = 176 kJ/mol, ΔS° = 0.285 kJ/(mol·K)

We note that the Gibbs free energy, ΔG° is represented by the following equation;

ΔG° = ΔH° - T·ΔS°

Where:

T = Temperature (Kelvin)

The reaction will be spontaneous for exergonic reactions, ΔG° < 0 and it will not be spontaneous for endergonic reaction, ΔG° > 0

At room temperature, T = 25 + 273.15 = 298.15 K

Which gives;

ΔG° = 176  - 298.15 × 0.285 = 91.03 kJ/mol which is > 0 Not spontaneous reaction

At 800°C, we have;

T = 273.15 + 800°C  + 1073.15 K

ΔG° = 176  - 1073.15 * 0.285 = -129.85 kJ/mol which is < 0 the reaction will be spontaneous

The correct option is therefore, that at room temperature, the reaction is not spontaneous. However, at high temperatures. like 800 °C, the free energy value turns negative and this reaction becomes spontaneous.

i dont know

explanation: