What is the fundamental difference between the effect of temperature and the effect of pressure on $K_p$?

Temperature changes the actual numerical value of $K_p$, whereas pressure only shifts the equilibrium position to maintain the existing $K_p$ value. Temperature is the only factor that can alter the equilibrium constant.

How does the response of an exothermic reaction to a temperature increase compare to that of an endothermic reaction?

An increase in temperature shifts an exothermic reaction to the left (decreasing $K_p$) to absorb heat, while it shifts an endothermic reaction to the right (increasing $K_p$) to consume the added thermal energy.

In a reaction where the number of moles of gaseous reactants equals the number of moles of gaseous products, how does a pressure increase affect the system?

A pressure increase has no effect on the equilibrium position if the moles of gas are equal on both sides. Since neither side reduces the total number of particles, the system cannot counteract the pressure change through a shift.

What is a common misconception regarding the effect of a catalyst on the yield of an equilibrium reaction?

A common error is believing a catalyst increases the yield of products. In reality, a catalyst only increases the rate at which equilibrium is reached; it does not change the final concentrations or the position of equilibrium.

Why is it incorrect to say that increasing the total pressure of a gaseous system always shifts the equilibrium to the right?

The shift depends entirely on the stoichiometry of the reaction. The equilibrium will shift toward whichever side has the fewer number of gas moles, which could be the reactants (left) or the products (right).

What mistake do students often make when counting moles for a pressure-related equilibrium shift involving solids and gases?

Students often include the moles of solids or liquids in their count. Only gaseous species are affected by pressure changes, so only the coefficients of substances with the $(g)$ state symbol should be summed.

Define Le Chatelier's Principle in the context of equilibrium.

It is the principle stating that if a system at equilibrium is disturbed by a change in temperature, pressure, or concentration, the system will shift its position to partially counteract the effect of that disturbance.

What does the term 'equilibrium position' refer to?

It refers to the relative proportions of reactants and products present at equilibrium. A shift 'to the right' means the proportion of products has increased relative to reactants.

What is an endothermic reaction in terms of enthalpy change?

An endothermic reaction is one that absorbs heat from the surroundings, characterized by a positive enthalpy change ($+\Delta H$). In these reactions, heat can be thought of as a reactant.

Why does $K_p$ decrease when temperature is increased for an exothermic reaction?

Increasing temperature adds heat, which the system counteracts by shifting in the endothermic (reverse) direction. This increases reactant partial pressures and decreases product partial pressures, leading to a smaller $K_p$ ratio.

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Advanced Physical Chemistry

Changes in Temperature & Pressure

Summary

This knowledge module explores how external changes in temperature and pressure influence gaseous equilibria. It distinguishes between shifts in the position of equilibrium (kinetic/spatial adjustments) and changes to the equilibrium constant $K_p$ (thermodynamic adjustments), while also clarifying the role of catalysts in these systems.

1. Le Chatelier's Principle in Gaseous Systems

Core Definition: Le Chatelier's Principle states that if a system at equilibrium is subjected to a change in conditions, the system will shift its equilibrium position to counteract that change. This principle provides a qualitative framework for predicting how a reaction mixture responds to external stressors.
Dynamic Response: When a stress is applied, the rates of the forward and reverse reactions temporarily become unequal. The system eventually reaches a new state of equilibrium where the rates are again equal, but the relative amounts of reactants and products may have changed.
Application Scope: In gaseous systems, this principle is applied specifically to changes in temperature, total pressure, and the addition of catalysts, each affecting the system through different physical mechanisms.

2. Temperature Effects and the Equilibrium Constant

3. Pressure Changes and Equilibrium Position

Diagram showing a piston compressing a gas mixture, illustrating how increased pressure forces the system to shift towards the side with fewer molecules to reduce the stress.

4. The Role of Catalysts

5. Key Distinctions: $K_p$ vs. Equilibrium Position

6. Exam Strategy & Tips