At 298K, Go = -67.4 kJ for the reaction
CO + Cl2 --> CoCl2
Calculate the change in Gibbs free energy (in kJ) at the same temperature when P (CO) = 0.368 bar, P (Cl2) = 0.62 bar and P (COCl2) = 0.872 bar. (R = 8.314 J/K mol)
Bump
Do you still need help on this one?
Yes.
Try this but check my thinking. Let me know how this turns out.
delta G = delta Go + RT ln Q.
I would change bars to atmospheres and use Q = PCOCl2/PCO*PCl2
To calculate the change in Gibbs free energy (ΔGo) at the given conditions, we will use the equation:
ΔGo = Go + RT ln(Q)
Where:
ΔGo = Change in Gibbs free energy
Go = Standard Gibbs free energy change at standard conditions
R = Gas constant (8.314 J/K mol)
T = Temperature (298 K)
ln = Natural logarithm
Q = Reaction Quotient
First, let's calculate the reaction quotient (Q) using the given partial pressures:
Q = (P(COCl2) / P(CO))^n * (P(Cl2))^m
Here, n and m are the stoichiometric coefficients of the balanced equation.
From the balanced equation:
CO + Cl2 --> CoCl2
We can see that the stoichiometric coefficient of CO is 1 (n = 1) and the stoichiometric coefficient of Cl2 is also 1 (m = 1).
Plugging in the given partial pressures:
Q = (0.872 / 0.368)^1 * (0.62)^1
Now, let's calculate Q:
Q = 2.37 * 0.62
Q ≈ 1.470
Now, we can substitute the values into the equation:
ΔGo = -67.4 kJ + (8.314 J/K mol * 298 K) * ln(1.470)
Let's calculate ΔGo:
ln(1.470) ≈ 0.388
ΔGo = -67.4 kJ + (8.314 J/K mol * 298 K) * 0.388
ΔGo ≈ -67.4 kJ + 971.411 kJ
ΔGo ≈ 904.011 kJ
Therefore, the change in Gibbs free energy (ΔGo) at the given conditions is approximately 904.011 kJ.