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Answer :
The Gibbs free energy change (ΔGº) for this reaction is approximately -38.778 kJ.Option A
To find the Gibbs free energy change (ΔGº), we need to use the equation ΔGº = -RTlnK where:
- R is the gas constant.
- T is the absolute temperature in Kelvin.
- K is the equilibrium constant.
- ln is the natural logarithm function, also known as the logarithm base e.
Given that the equilibrium constant (K) is 6.32 * 10⁶, the temperature (T) is 298 K, and the gas constant (R) is 8.31 * 10⁻³ kJ/K.
Firstly, we start by multiplying the temperature (T) by the gas constant (R). That is 298 K * 8.31 * 10⁻³ kJ/K.
Next, we take the natural logarithm (ln) of the equilibrium constant (K). That is ln(6.32 * 10⁶).
We then multiply these two results together. Hence the equation becomes. -RTlnK.
Remember that ΔGº shows the direction and magnitude of how a chemical reaction will proceed. If the ΔGº is negative, this means that the reaction will release energy and hence it will occur spontaneously. The larger the absolute value, the more likely the reaction is to occur.
By substituting the given numbers into the equation, we calculate and find that the Gibbs free energy change (ΔGº) for this reaction is approximately -38.778 kJ.
So the answer to the problem is -38.778 kJ. Therefore, the reaction will proceed spontaneously.
Option A
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