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A negative ΔG, or Gibbs free energy change, indicates that the reaction is thermodynamically favored to occur spontaneously under the given conditions. This means that the products of the reaction have lower free energy than the reactants, allowing the reaction to proceed without needing additional external energy input. In this context, spontaneous does not necessarily mean that the reaction happens rapidly; rather, it signifies a tendency toward the formation of products from the reactants.

In a scenario where ΔG is negative, the reaction will proceed in the direction that decreases the system's free energy, aligning with the principles of thermodynamics. This is crucial for understanding how energy flows in chemical processes and predicting the feasibility of reactions based on their Gibbs free energy.

When ΔG is equal to zero, the system is at equilibrium, meaning there is no net change in the concentration of reactants and products. A positive ΔG would signify that a reaction is nonspontaneous, suggesting that external energy would be required to drive it forward. Thus, recognizing the implications of a negative ΔG helps in predicting the behavior of chemical reactions and their favorability.