Identify the correct statement for the following reaction:

Br2 (l) → Br2 (g) ΔH° = 30.9 kJ ΔS° = 93.2 J/K


A) The reaction is always spontaneous.

B) The reaction is never spontaneous.

C) The reaction is at equilibrium; therefore, ΔG° = 0.

D) The reaction could be spontaneous at high temperatures.

E) The reaction could be spontaneous at low temperature.

Used the following equation and plug in a couple of variables for T and see what you get.

ΔG=ΔH−TΔS

Remember, if ΔG is negative then the reaction is spontaneous, if positive, than the reaction is not spontaneous, and if equal to 0 is at equilibrium.

To determine the spontaneity of the reaction, we can use the Gibbs free energy equation:

ΔG° = ΔH° - TΔS°

where ΔG° is the standard Gibbs free energy change, ΔH° is the standard enthalpy change, ΔS° is the standard entropy change, and T is the temperature in Kelvin.

In this case, we have ΔH° = 30.9 kJ and ΔS° = 93.2 J/K. If the temperature is not provided, we can assume it is at a reasonable temperature, such as room temperature (298 K).

Plugging the values into the equation, we get:

ΔG° = 30.9 kJ - (298 K) * (93.2 J/K) = 30.9 kJ - 27.8 kJ = 3.1 kJ

Since ΔG° is positive, this means the reaction is non-spontaneous.

The correct statement for the reaction Br2 (l) → Br2 (g) is:

B) The reaction is never spontaneous.

To determine the spontaneity of a reaction, we use the concept of Gibbs free energy (ΔG). The relationship between ΔG, ΔH (enthalpy change), and ΔS (entropy change) is given by the equation:

ΔG = ΔH - TΔS

Where T is the temperature in Kelvin.

Given that ΔH° = 30.9 kJ and ΔS° = 93.2 J/K for the reaction Br2 (l) → Br2 (g), we can determine the spontaneity of the reaction by analyzing the sign of ΔG.

If ΔG < 0, the reaction is spontaneous.
If ΔG > 0, the reaction is non-spontaneous.
If ΔG = 0, the reaction is at equilibrium.

Now, let's substitute the given values into the equation and analyze the sign of ΔG:

ΔG = 30.9 kJ - T(93.2 J/K)

Since the units are different, we need to convert kJ to J by multiplying by 1000:

ΔG = 30,900 J - T(93.2 J/K)

From the equation, we can see that ΔG will be negative if ΔH is greater than TΔS.

Let's consider the possible scenarios based on temperature:

1) At low temperature (T), TΔS will be small compared to 30,900 J. In this case, ΔG will likely be positive, indicating that the reaction is non-spontaneous. Therefore, option E) The reaction could be spontaneous at low temperature is incorrect.

2) At high temperature (T), TΔS will be large compared to 30,900 J. In this case, ΔG will likely be negative, indicating that the reaction is spontaneous. Therefore, option D) The reaction could be spontaneous at high temperatures is correct.

Hence, the correct statement for the given reaction is:

D) The reaction could be spontaneous at high temperatures.