calculate the equilibrium constant for the following reaction at 298k given that DeltaGrxn=2.8kj/mol N2O4-2NO2
I thought to get K you use the expression K=e^(-deltaG/(RT)) but I cant get the same answer my professor put down which is 3.1
I get his answer. Did you use J, or kJ in the exponent? kJ is not a standard unit, try 2800J/mol
PCl5(g) eqarrow PCl3(g) + Cl2(g)
The reaction above has an equilibrium constant of 0.800 at 340°C, and deltaH°rxn = 87.9 kJ/mol under standard conditions. Under which conditions will this reaction produce the most PCl3(g)?
To calculate the equilibrium constant (K) for a reaction given the standard Gibbs free energy change (ΔG°), temperature (T), and the balanced equation, you can use the following expression:
K = e^(-ΔG° / (RT))
Where:
- K is the equilibrium constant.
- ΔG° is the standard Gibbs free energy change (in this case, 2.8 kJ/mol).
- R is the ideal gas constant (8.314 J/(mol·K)).
- T is the temperature in Kelvin (298 K).
However, in your calculation, you converted the units of ΔG° from kJ to J, which led to a different answer. It is important to keep consistent units throughout the calculation. Therefore, let's convert ΔG° to J:
ΔG° = 2.8 kJ/mol = 2.8 × 10^3 J/mol
Now, substitute the values into the equation:
K = e^(-ΔG° / (RT))
= e^(-(2.8 × 10^3 J/mol) / (8.314 J/(mol·K) × 298 K))
Calculating this expression will give you the equilibrium constant (K) for the reaction at 298 K.