At 298 K one reaction has K = 1.0 × 10³ and another has K = 1.0 × 10⁻²…

For $a\text{A} + b\text{B} \rightleftharpoons c\text{C} + d\text{D}$: $K_c = \dfrac{[\text{C}]^c[\text{D}]^d}{[\text{A}]^a[\text{B}]^b}$. $K_p = K_c(RT)^{\Delta n}$ with $R = 0.0831$ L bar K$^{-1}$ mol$^{-1}$. Pure solids and liquids are left out. $Q_c < K_c$: forward; $Q_c > K_c$: backward. $\Delta G^\circ = -2.303\,RT\log K$.