$\overline{){\mathbf{\u2206}}{\mathbf{G}}{{\mathbf{\xb0}}}_{{\mathbf{rxn}}}{\mathbf{}}{\mathbf{=}}{\mathbf{}}{\mathbf{-}}{\mathbf{RT}}{\mathbf{}}{\mathbf{lnK}}}\phantom{\rule{0ex}{0ex}}\mathbf{ln}\mathbf{}\mathbf{K}\mathbf{}\mathbf{=}\mathbf{}\frac{\mathbf{\u2206}\mathbf{G}{\mathbf{\xb0}}_{\mathbf{rxn}}}{\mathbf{-}\mathbf{RT}}\phantom{\rule{0ex}{0ex}}\overline{){\mathbf{K}}{\mathbf{}}{\mathbf{=}}{\mathbf{}}{{\mathbf{e}}}^{\frac{\mathbf{\u2206}\mathbf{G}{\mathbf{\xb0}}_{\mathbf{rxn}}}{\mathbf{-}\mathbf{RT}}}}$

The following equation represents the decomposition of a generic diatomic element in its standard state.

1/2x^2 (g) yields X(g)

assume that the standard molar Gibbs energy of formation of X(g) is 4.45 KJ *mol^-1 at 2000. K and -51.20 KJ *mol^-1 at 3000. K. Determine the value of K (the thermodynamics equilibrium constant) at each temperature.

At 2000. K, we were given: delta Gf =4.45 kj*mol^-1. what is K at the temperature?

K at 2000. K = ?

At 3000. K, we are given: delta Gf =-51.20 KJ* mol^-1. what is the K at that temperature?

K at 3000. K= ?

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