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Chemical equilibrium occurs when the rate of the forward reaction equals the rate of the reverse reaction. At this point, the concentrations of reactants and products remain constant over time. It is important to understand that even though the concentrations do not change, the reactions are still occurring. This dynamic balance ensures that the system is in a state of equilibrium. For the reaction, \[ N_2(g) + 3 H_2(g) \rightleftharpoons 2 NH_3(g) \], the equilibrium state is reached when the formation of ammonia (NH_3) from nitrogen (N_2) and hydrogen (H_2) is balanced by the decomposition of ammonia back into nitrogen and hydrogen. Remember, equilibrium can be reached from both directions: whether starting with reactants or products.

Gas reactions involve reactants and products in the gaseous state. The concentration of gases is often expressed in molarity (moles per liter) or partial pressure. In our example, the reaction of nitrogen (N_2) and hydrogen (H_2) to form ammonia (NH_3) is a gas-phase reaction. Understanding gas reactions involves being familiar with how volume and pressure changes can affect gas concentrations. Moreover, gases generally diffuse and intermingle at faster rates compared to solids and liquids because of the higher energy and movement of gas molecules.

The reaction quotient, Q, compares the concentration of products to reactants at any point in the reaction, not just at equilibrium. It is expressed as: \[Q = \frac{[Products]}{[Reactants]}\text{ raised to their respective stoichiometric coefficients.} \] For our given reaction, it’s represented as: \[Q = \frac{[NH_3]^2}{[N_2][H_2]^3}\text{.} \] If Q is less than the equilibrium constant (K), the reaction will proceed forward to produce more products. If Q is greater than K, the reaction will shift backward to produce more reactants. If Q equals K, the system is at equilibrium. In the exercise, calculating Q with the equilibrium concentrations confirms the reaction’s closeness to K, providing insights into which direction the reaction might shift.

In chemical reactions, the change in concentration involves understanding how the quantities of reactants and products vary over time before reaching equilibrium. For instance, in the given reaction, we start with initial concentrations: \[ [N_2] = 1 M \] \[ [H_2] = 3 M \] \[ [NH_3] = 0 M \] When the system reaches equilibrium, changes occur: \[ [NH_3] = 0.05 M \] The change in concentration of NH_3 provides insight into the consumption of reactants: \[ \text{Change in } [N_2] = -\frac{0.05 M}{2} = -0.025 M \] \[ \text{Change in } [H_2] = -3 \times \frac{0.05 M}{2} = -0.075 M \] Thus, the equilibrium concentrations become: \[ [N_2]_{eq} = 1 - 0.025 = 0.975 M \] \[ [H_2]_{eq} = 3 - 0.075 = 2.925 M \] These changes help us understand how the reaction progresses towards equilibrium.