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Chemical equilibrium occurs in a reversible reaction when the rates of the forward and reverse reactions are equal. This means that the concentrations of all reactants and products remain constant over time. It is important to understand that equilibrium is dynamic, not static, meaning the reactions continue to occur, but there is no net change in concentration.

In chemical equilibrium, the system will spontaneously shift to counteract any changes made to the system, as described by Le Chatelier’s Principle. If you add more reactant or remove product, the system will adjust to minimize the change by favoring the forward reaction. Conversely, adding product or removing reactant will shift the equilibrium to favor the reverse reaction.

Understanding chemical equilibrium is crucial for predicting how different factors like concentration, temperature, and pressure affect the reaction. By controlling these variables, chemists can manipulate the conditions to favor the desired reaction direction.

Equilibrium expressions, or equilibrium constant expressions, are mathematical representations of a reaction at equilibrium. They allow chemists to determine the ratio of product concentrations to reactant concentrations.

For reactions where all reactants and products are gases or solutions, the equilibrium expression for the reaction \( aA + bB \rightleftharpoons cC + dD \) is given by \( K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b} \).

Note that in equilibrium expressions, the concentrations of solids and pure liquids are omitted as their activities are constant and taken as 1. Only the concentrations of aqueous and gaseous species are included. This is why in reactions (a) through (d), only gaseous or aqueous species contribute to the equilibrium constant \( K_c \).

• Solids and liquids are not included in the expression.
• Only consider aqueous solutions and gases.
• The coefficients of the reaction become the exponents in the expression.

Knowing how to construct equilibrium expressions allows us to predict the behavior of a system in equilibrium, especially how it will respond to changes.

Gaseous reactions are chemical reactions where the reactants and products are in the gas phase. An interesting aspect of such reactions is how they respond to changes in pressure and temperature.

In gaseous equilibria, changes in pressure can disturb the equilibrium. According to Le Chatelier's Principle, if you increase the pressure by decreasing the volume, the equilibrium will shift towards the side with the fewer gas molecules. Conversely, decreasing the pressure by increasing the volume favors the side with more gas molecules.

• In reaction \((a)\), the effect of pressure change would alter the concentrations of \( \mathrm{NH}_3 \) and \( \mathrm{HCl} \).
• In reaction \((b)\), increasing pressure could favor forming \( \mathrm{CO}_2 \) and \( \mathrm{N}_2 \) gases, as the forward reaction consumes more gas molecules than it produces.

Understanding how gaseous reactions behave can help us harness this knowledge in industrial applications, such as optimizing yields in chemical manufacturing processes by manipulating reaction conditions.