What is the fundamental difference between a reaction that goes to completion and a reversible reaction?

A reaction that goes to completion proceeds until at least one reactant is entirely consumed. In contrast, a reversible reaction allows products to react back into reactants, eventually reaching a state where both processes occur at the same rate.

How does a closed system differ from an open system in the context of chemical equilibrium?

A closed system allows for the exchange of energy but not matter with its surroundings, which is necessary for equilibrium to be established. An open system allows matter to escape, which prevents the backward reaction from balancing the forward reaction.

What is the difference between the 'rate of reaction' and the 'position of equilibrium' when a catalyst is added?

A catalyst increases the rate of reaction by allowing equilibrium to be reached faster. However, it has no effect on the position of equilibrium because it speeds up the forward and backward reactions by the exact same factor.

What is a common error when writing the $K_c$ expression for a reaction involving a solid reactant?

A common mistake is including the solid in the expression. Solids must be excluded because their concentration is constant and does not change the ratio of the equilibrium constant.

Why is it incorrect to say that $K_c$ increases when the concentration of a reactant is increased?

Increasing reactant concentration shifts the position of equilibrium to the right, but the ratio of products to reactants at the new equilibrium will still equal the original $K_c$ value. Only temperature can change the value of $K_c$.

What mistake do students often make when applying pressure changes to a reaction with equal moles of gas on both sides?

Students often assume a shift will occur, but if the number of gas moles is equal on both sides, a change in pressure has no effect on the position of equilibrium. The system cannot counteract the pressure change by shifting.

Define 'Dynamic Equilibrium' in terms of reaction rates and concentrations.

Dynamic equilibrium is a state where the rate of the forward reaction equals the rate of the backward reaction. Consequently, the concentrations of all reactants and products remain constant over time.

State Le Chatelier's Principle.

Le Chatelier's Principle states that if a system at equilibrium is disturbed by a change in conditions, the position of equilibrium will shift in a direction that tends to counteract or oppose that change.

What does the magnitude of $K_c$ indicate about the extent of a reaction?

A very large $K_c$ (much greater than 1) indicates that the equilibrium position lies far to the right, favoring products. A very small $K_c$ (much less than 1) indicates that the equilibrium lies far to the left, favoring reactants.

Why does increasing the temperature of an exothermic reaction decrease the value of $K_c$?

In an exothermic reaction, heat is treated as a product. Increasing temperature shifts the equilibrium to the left (the endothermic direction) to absorb heat, decreasing product concentration and increasing reactant concentration, which lowers $K_c$.

Chemical Equilibria | AQA A-Level Chemistry

1. Definition & Core Concepts

Reversible Reactions: Unlike reactions that go to completion, reversible reactions allow products to react together to reform the original reactants. This is denoted by the equilibrium arrow ( $\rightleftharpoons$ ), signifying that both forward and backward processes are occurring simultaneously.
Dynamic Equilibrium: This state is reached when the rate of the forward reaction equals the rate of the backward reaction in a closed system. While the concentrations of reactants and products remain constant, the system is 'dynamic' because molecules are still reacting at the microscopic level.
Closed vs. Open Systems: A closed system prevents the exchange of matter with the surroundings, which is a prerequisite for establishing equilibrium. In an open system, products (especially gases) can escape, driving the reaction toward completion and preventing the backward reaction from balancing the forward one.

A graph showing concentration vs time for a reversible reaction. Reactant concentration decreases and product concentration increases until both become constant at the point where equilibrium is reached.

2. Underlying Principles: Le Chatelier's Principle

The Principle: If a system at equilibrium is subjected to a change in conditions (concentration, pressure, or temperature), the system will shift its position of equilibrium to counteract that change. This is a qualitative tool used to predict how a system responds to external stress.
Concentration Effects: Increasing the concentration of a reactant causes the system to shift to the right (product side) to consume the excess. Conversely, removing a product will also shift the equilibrium to the right to replace what was lost.
Pressure Effects: Pressure changes only affect systems involving gases where there is a difference in the number of moles of gas between reactants and products. Increasing pressure shifts the equilibrium toward the side with the fewer number of gas molecules to reduce the total pressure.

3. Temperature and Catalysts

Temperature Effects: The direction of the shift depends on whether the forward reaction is exothermic (releases heat) or endothermic (absorbs heat). Increasing temperature favors the endothermic direction to absorb the added thermal energy, while decreasing temperature favors the exothermic direction.
Catalyst Role: A catalyst increases the rate of both the forward and backward reactions equally by providing an alternative pathway with lower activation energy. Consequently, a catalyst allows a system to reach equilibrium faster but does not change the position of equilibrium or the yield of products.

4. Methods: The Equilibrium Constant ($K_c$)

The Expression: For a general reaction $aA + bB \rightleftharpoons cC + dD$ , the equilibrium constant is defined as $K_c = \frac{[C]^c [D]^d}{[A]^a [B]^b}$ . The square brackets denote concentrations in $\text{mol dm}^{-3}$ at equilibrium.
Exclusions: Pure solids and liquids are excluded from the $K_c$ expression because their concentrations are effectively constant. Only aqueous species and gases are included in the calculation.
Units of $K_c$ : The units are not fixed and must be derived for every specific reaction by substituting the units of concentration into the expression and canceling them out. If the number of moles on both sides is equal, $K_c$ will be unitless.

5. Key Distinctions

Feature	Position of Equilibrium	Equilibrium Constant ( $K_c$ )
Definition	Relative amounts of reactants and products present.	A numerical value derived from the ratio of concentrations.
Affected by Concentration	Yes, shifts to counteract change.	No, remains constant.
Affected by Pressure	Yes (if gas moles differ).	No, remains constant.
Affected by Temperature	Yes, shifts based on enthalpy.	Yes, the value of $K_c$ changes.
Affected by Catalyst	No effect.	No effect.

Yield vs. Rate: In industrial processes, conditions are often a 'compromise.' For example, a low temperature might give a high equilibrium yield but a reaction rate that is too slow for commercial viability. Therefore, intermediate temperatures are used to balance yield and speed.

6. Exam Strategy & Tips

Temperature is Unique: Always remember that temperature is the only factor that changes the numerical value of $K_c$ . If an exam question asks about the effect of concentration or pressure on $K_c$ , the answer is always 'no change.'
ICE Tables: When calculating $K_c$ from initial amounts, use an Initial, Change, Equilibrium table. Use the stoichiometric ratios from the balanced equation to determine the 'Change' row based on the known equilibrium amount of one species.
Check the State Symbols: Before writing a $K_c$ expression, check for $(s)$ or $(l)$ symbols. Including a solid in your expression is a common mistake that will result in the loss of marks.
Units Derivation: Never assume units. Always write out the units for each term in the expression (e.g., $\frac{(\text{mol dm}^{-3})^2}{(\text{mol dm}^{-3})^3}$ ) and simplify carefully.

Chemical Equilibria

Summary

1. Definition & Core Concepts

2. Underlying Principles: Le Chatelier's Principle

3. Temperature and Catalysts

4. Methods: The Equilibrium Constant ($K_c$)

5. Key Distinctions

6. Exam Strategy & Tips

Chemical Equilibria

Summary

1. Definition & Core Concepts

2. Underlying Principles: Le Chatelier's Principle

3. Temperature and Catalysts

4. Methods: The Equilibrium Constant ($K_c$)

5. Key Distinctions

6. Exam Strategy & Tips