What is the fundamental difference between a static equilibrium and a dynamic chemical equilibrium?

In static equilibrium, all motion and change have stopped. In dynamic chemical equilibrium, the forward and reverse reactions continue to occur at equal rates, so while concentrations remain constant, the molecules are still actively reacting.

How does the use of a catalyst affect the value of the equilibrium constant $K$?

A catalyst has no effect on the value of $K$. It increases the rate of both the forward and reverse reactions equally, allowing the system to reach equilibrium faster without changing the final ratio of products to reactants.

When comparing $Q$ and $K$, what does it mean if $Q > K$?

If $Q > K$, the current ratio of products to reactants is higher than the equilibrium ratio. The system will shift in the reverse direction (toward the reactants) to reach equilibrium.

What is a common mistake when writing the equilibrium expression for a reaction involving a solid reactant?

A common error is including the solid in the $K_c$ expression. Pure solids and liquids must be excluded because their concentrations do not change significantly during the reaction.

Why is it incorrect to say that 'the reaction stops' once equilibrium is reached?

The reaction is dynamic; reactants continue to form products and products continue to reform reactants. The 'stop' is only apparent at the macroscopic level because the rates are perfectly balanced.

What happens if you apply Le Chatelier's Principle to a pressure change in a system where the moles of gas are equal on both sides?

The system will not shift in either direction. Pressure changes only affect equilibrium positions when there is a difference in the number of moles of gaseous reactants and products.

Define the Law of Mass Action.

The Law of Mass Action states that for a reversible chemical reaction at equilibrium at a constant temperature, a certain ratio of reactant and product concentrations has a constant value ($K$).

What is the formula for $K_p$ in terms of partial pressures?

For a reaction $aA + bB \rightleftharpoons cC + dD$, $K_p = \frac{(P_C)^c (P_D)^d}{(P_A)^a (P_B)^b}$, where $P$ represents the partial pressure of each gaseous component at equilibrium.

How is the Reaction Quotient $Q$ mathematically related to the Equilibrium Constant $K$?

Both $Q$ and $K$ use the same stoichiometric expression (products over reactants raised to their coefficients). The difference is that $K$ uses equilibrium concentrations, while $Q$ uses concentrations at any specific time.

Why does increasing the temperature of an endothermic reaction increase the value of $K$?

In an endothermic reaction, heat acts like a reactant. Adding heat (increasing temperature) drives the reaction forward to consume the energy, resulting in a higher concentration of products and thus a larger $K$.

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Reactions & Equilibrium

Summary

Chemical equilibrium is a dynamic state in reversible reactions where the rates of the forward and reverse processes are equal, resulting in constant concentrations of reactants and products over time. This state is governed by the Law of Mass Action and can be shifted by external changes in concentration, pressure, or temperature according to Le Chatelier's Principle.

1. Nature of Reversibility and Dynamic Equilibrium

Reversible reactions are chemical processes that can proceed in both the forward (reactants to products) and reverse (products to reactants) directions simultaneously. Unlike irreversible reactions that go to completion, reversible reactions reach a state where both reactants and products coexist in the system.
Dynamic equilibrium occurs when the rate of the forward reaction exactly equals the rate of the reverse reaction. Although the macroscopic properties like concentration, pressure, and color remain constant, the microscopic processes continue to occur at equal rates.
A closed system is a prerequisite for establishing equilibrium in reactions involving gases. This ensures that no matter or energy (in the form of mass) escapes, allowing the products to collide and reform the reactants.

Concentration vs Time graph showing reactants decreasing and products increasing until they reach constant levels at equilibrium.

2. The Equilibrium Constant ($K$)

3. Reaction Quotient ($Q$) and Directionality

4. Le Chatelier's Principle

Le Chatelier's Principle states that if a system at equilibrium is disturbed by a change in conditions, the system will shift its equilibrium position in a direction that tends to counteract the disturbance.
Concentration Changes: Increasing the concentration of a reactant shifts the equilibrium toward the products to consume the added substance. Conversely, removing a product shifts the equilibrium toward the products to replace what was lost.
Pressure and Volume: For gaseous systems, decreasing the volume (increasing pressure) shifts the equilibrium toward the side with the fewer moles of gas. Increasing the volume shifts it toward the side with more moles of gas.
Temperature: This is the only factor that changes the actual value of $K$ . For exothermic reactions, increasing temperature shifts the equilibrium toward reactants ( $K$ decreases); for endothermic reactions, it shifts toward products ( $K$ increases).

5. Key Distinctions

6. Exam Strategy & Tips