Final answer:
When reacting 4.0 mol of A with 10.0 mol of B in a 1:1 molar ratio, A is the limiting reactant. Adding 0.50 mol of A to an equilibrium system will shift the reaction to produce more B. Equilibrium adjustments and stoichiometric relationships are analyzed using ICE tables and equilibrium constants.
Step-by-step explanation:
Based on the provided snippets of information that discuss reactions involving substances A and B, the addition of moles of reactants, and the concept of chemical equilibrium, we are dealing with a stoichiometry problem that addresses the Law of Definite Proportions and the Law of Conservation of Mass. When 4.0 mol of A is reacted with 10.0 mol B, we must consider the mole ratios provided by the balanced chemical equation to determine the limiting reactant and the extent of the reaction.
If the mole ratio is 1:1, as it is in some of the given cases, then A would be the limiting reactant, and the reaction would produce 4.0 mol B according to the stoichiometry of the equation. Additionally, when a substance is added to a system at equilibrium, the system will shift according to Le Châtelier's Principle. In this context, adding 0.50 moles of A would cause the equilibrium to shift to the right, resulting in the production of more B.
In terms of equilibrium calculations, ICE tables are used to visualize initial concentrations, changes in concentrations due to reactions, and new equilibrium concentrations after the system has adjusted. When we talk about changes in gas moles (Δn) and temperature (T), we are usually considering the application of the ideal gas law and equilibrium constants (K) to predict how the system will respond to changes in conditions, following the equilibrium constant expression that generally takes on the form K = [products]/[reactants], raised to the power of their stoichiometric coefficients in the balanced equation.