Calculate the temperature at which it is thermodynamically possible for C (s) to reduce iron(III) oxide by the reaction:
2Fe2O3 (s) + 3 C (s) ---> 4Fe (s) + 3CO2
I don't need a specific answer, but I'm not sure how to approach the problem.
I'm assuming it has to do with substituting values in of the reaction into the change in free energy equals the change in enthalpy minus (temperature * change in entropy). Would it be the temperature when the reaction is negative AKA spontaneous
To determine the temperature at which it is thermodynamically possible for C (s) to reduce iron(III) oxide (Fe2O3), we can use the concept of Gibbs free energy (ΔG).
First, we need to know the standard Gibbs free energy change (ΔG°) for the reaction. This value can be found in thermodynamic tables or calculated using the standard enthalpy change (ΔH°) and standard entropy change (ΔS°):
ΔG° = ΔH° - T * ΔS°
Where:
ΔG° is the standard Gibbs free energy change
ΔH° is the standard enthalpy change
T is the temperature in Kelvin
ΔS° is the standard entropy change
To find the temperature at which the reaction is thermodynamically possible (i.e., when ΔG° is negative), we can rearrange the equation:
T = (ΔH° / ΔS°)
Here, the temperature will be in Kelvin (K).
To solve this, you will need the values of ΔH° and ΔS° for the given reaction. These values can be found in thermodynamic databases or textbooks. Once you have those values, substitute them into the equation above, and you will get the temperature at which the reaction is thermodynamically feasible.
Keep in mind that the reaction you provided is the balanced equation and assumes standard conditions (1 atm pressure and 298 K temperature). Depending on the actual conditions, the reaction might require different stoichiometric ratios or may not proceed at all.
To determine the temperature at which it is thermodynamically possible for C (s) to reduce iron(III) oxide, we can use the Gibbs Free Energy equation:
ΔG = ΔH - TΔS
In this equation, ΔG represents the change in free energy, ΔH represents the change in enthalpy, T represents the temperature in Kelvin, and ΔS represents the change in entropy.
If the reaction is spontaneous (i.e., thermodynamically possible), then ΔG will be negative. Therefore, we can rewrite the equation as:
ΔG < 0
ΔH - TΔS < 0
Solving this equation for temperature (T), we get:
T > ΔH/ΔS
So the temperature at which it is thermodynamically possible for C (s) to reduce iron(III) oxide is greater than the ratio of the change in enthalpy (ΔH) to the change in entropy (ΔS) for the reaction.
To obtain more specific numerical values, you would need to know the actual values of ΔH and ΔS for the reaction and substitute them into the equation.