Final answer:
A chemical reaction is at equilibrium when the forward and reverse reactions occur at equal rates, and the concentrations of products and reactants are constant but not necessarily equal. A reaction is not at equilibrium if the concentration of reactants is increasing or the forward reaction is very slow. Equal concentrations of reactants and products do not definitively indicate whether a system is at equilibrium.
Step-by-step explanation:
Conditions that indicate a chemical reaction is at equilibrium include: the forward and reverse reactions happening at equal rates, and the concentrations of products and reactants being constant. However, these concentrations do not have to be equal, they just don't change over time. For example, in the equilibrium A+B=C+D, both the forward (A+B to C+D) and the reverse (C+D to A+B) reactions are occurring at the same rate, making the reaction dynamic.
A reaction is not at equilibrium if the concentration of reactants is slowly increasing, or if the forward reaction is proceeding at a very slow rate. This indicates that the system is still undergoing change and has not yet reached a state where the rates of the forward and reverse reactions are balanced.
Scenarios that may or may not indicate equilibrium are when the products and reactants have equal concentrations, or when the concentration of products is greater than that of reactants. These conditions alone do not determine whether the system is at equilibrium, as equilibrium is defined by the rates of the reactions and not solely the concentration of reactants and products.