Calculate the temp at which the reaction (2POCL3(g) -> 2PCl3(g)+O2(g))would become spontaneous. Given: DH= +572kJ and DS= 179 J/K.

Would the answer be 298 K? I found DG to be 518.658 kJ, and then plugged in all the values to DG= DH-TDS.
Or, should I ignore the DG, set DH-TDS to = 0, and find the answer to be about 3195.53 K?
Which is the proper way to go about answering this question?

I think you do the latter. Set DG = 0 (or -1 since, technically, DG = 0 at equilibrium). I think these are handled by setting DG = 0, then stick an appendix to the answer that says "anything above this T."

To determine the temperature at which the reaction becomes spontaneous, you need to calculate the Gibbs free energy change (ΔG) and set it to zero. The proper way to solve this problem is by using the equation ΔG = ΔH - TΔS, where ΔH is the enthalpy change, ΔS is the entropy change, T is the temperature (in Kelvin), and ΔG is the Gibbs free energy change.

Given information:
ΔH = +572 kJ
ΔS = 179 J/K

First, convert ΔH to J:
ΔH = +572 kJ = +572,000 J

Next, convert ΔS from J/K to kJ/K:
ΔS = 179 J/K = 0.179 kJ/K

Now, we can substitute the values into the equation and solve for T:
ΔG = ΔH - TΔS

0 = 572,000 J - T * 0.179 kJ/K

Rearrange the equation to solve for T:
T * 0.179 kJ/K = 572,000 J

T = (572,000 J) / (0.179 kJ/K)
T ≈ 3195.53 K

Therefore, the correct answer is approximately 3195.53 K.

To determine the temperature at which a reaction becomes spontaneous, you need to use the Gibbs free energy equation: ΔG = ΔH - TΔS. In this equation, ΔG represents the change in Gibbs free energy, ΔH represents the change in enthalpy, ΔS represents the change in entropy, and T represents the temperature in Kelvin.

To find the temperature at which the reaction becomes spontaneous, you want to find the point where ΔG is equal to zero, as this indicates the transition from non-spontaneous to spontaneous.

In your first attempt, you correctly calculated ΔG to be 518.658 kJ. However, instead of plugging it into the equation DG = DH - TDS, you should set ΔG equal to zero and solve for T.

0 = ΔH - TΔS

Rearranging the equation:
TΔS = ΔH

Now, you can substitute the given values:

T * (179 J/K) = +572,000 J (since 572 kJ is equal to 572,000 J)

Now, convert both sides to the same units:

T * (179 J/K) = 572,000 J

Now divide both sides by 179 J/K:

T ≈ 3195.53 K

Therefore, the proper way to calculate the temperature at which the reaction would become spontaneous is approximately 3195.53 K, not 298 K.

Remember to verify your calculations and units during the process to ensure accuracy.