At room temperature (25 °C), the gas ClNO is impure because it decomposes slightly according to the equation
2ClNO(g) Cl2(g) + 2NO(g)
The extent of decomposition is about 4%. Based upon this, what is the approximate value of ΔG°298 for the reaction at that temperature?
Is my answer correct?:
Kp = (0,02* 0.04^2)/ (0.96^2)= 3.472E5
ΔG°298 = [-(8.314)*298*In(3.472E5)]1000J
ΔG°298 = 25.4J
I think the sign for exponent of Kp is -5 and not +5. Check that out.
I think you have too many significant figures in 3.472 but you have rounded correctly in the answer. .
I don't understand the dG calculation at all. If I plug in 3.472E+5 I don't get 25.4. If I plug in 3.472E-5 (Kp that should be used) I get 25.4 kJ and not 25.4 J.
To determine the approximate value of ΔG°298 for the reaction, you first need to calculate the equilibrium constant (Kp) at the given temperature. Here's the correct way to calculate it:
1. Start by representing the given extent of decomposition (4%) as a decimal: 4% = 0.04.
2. Since the reaction involves a decomposition of ClNO, the equilibrium constant expression can be written as:
Kp = [Cl2] * [NO]^2 / [ClNO]^2.
3. Using the stoichiometric coefficients from the balanced equation, write the equilibrium concentration expression as:
Kp = [Cl2] / [ClNO]².
4. Substitute the values into the equation:
Kp = (0.02) / (0.96^2) ≈ 0.0216.
Now, to calculate ΔG°298 using the obtained Kp value, follow these steps:
1. Use R = 8.314 J/(mol⋅K) as the ideal gas constant and T = 298 K as the temperature.
2. The Gibbs free energy change (ΔG°298) can be calculated using the equation:
ΔG°298 = -RT ln(Kp).
3. Plug in the values:
ΔG°298 = - (8.314 J/(mol⋅K)) * (298 K) * ln(0.0216).
ΔG°298 ≈ - 3478 J/mol.
Therefore, your answer of ΔG°298 = -3478 J/mol is correct, not 25.4 J as you suggested in your initial calculation.