CHEM 1036: Chapter 17

Chapter 17: Equilibrium - The Extent of Chemical Reactions

Introduction

• Instructor: Dr. Shamindri M. Arachchige

• Email: arachsm@vt.edu

Equilibrium Concepts

• Definition of Chemical Equilibrium: A state in which the concentrations of reactants and products remain constant over time while chemical reactions are still occurring in both directions.

• Relationship Between Equilibrium and Reaction Rates: At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction.

• Equilibrium Constant Expression (K_eq):

  • Expressed as the ratio of the concentrations (or partial pressures) of products to reactants, with each concentration raised to the power of its coefficient in the balanced equation.

  • Example: For a generic reaction aA + bB ⇄ cC + dD,[ K_{eq} = \frac{[C]^c[D]^d}{[A]^a[B]^b} ]

• Reaction Quotient (Q): The ratio of product concentrations to reactant concentrations at any point in time, used to predict the direction of the reaction.

Le Chatelier’s Principle

• Describes how a system at equilibrium responds to external changes or disturbances (i.e., changes in temperature, pressure, or concentration).

• The system will shift to counteract the disturbance and maintain equilibrium.

  • Disturbances include:

    • Changing the concentration of reactants/products.

    • Changing the pressure of the system.

    • Changing the temperature of the reaction.

Types of Reactions

• Complete Reactions: Reactions that go to completion with no reactants remaining.

  • Example: 2 Mg(s) + O2(g) → 2 MgO(s)

• Reversible Reactions: Reactions where both reactants and products are present at equilibrium.

  • Example: N2O4(g) ⇄ 2NO2(g)

Dynamic Nature of Equilibrium

• At equilibrium, concentrations of reactants and products do not change over time, but both forward and reverse reactions continue to occur at equal rates.

Example: Equilibrium in N2O4 and NO2 Reaction

• At equilibrium state:

  • Rate of N2O4 consumption = Rate of NO2 production.

Equilibrium Constant Determination

• Conduct experiments under varying initial conditions to determine K_eq for the reactions:

  • Experiment 1: Start with 1.00 atm of N2O4, results in 0.22 atm N2O4 and 1.56 atm NO2 at equilibrium.

  • Experiment 2: Start with 1.00 atm of NO2, results in 0.07 atm N2O4 and 0.86 atm NO2 at equilibrium.

  • Experiment 3: Start with both reactants, results in 0.42 atm N2O4 and 2.16 atm NO2 at equilibrium.

Reaction Quotient and Equilibrium Constant

• Q changes throughout the reaction:

  • At equilibrium, Q = K

• K remains constant at a given temperature.

Pressure and Concentration Relationships

• K_p: Based on partial pressures of gases, typically denoted in atmospheres.

• K_c: Based on molar concentrations of reactants and products.

• Solids and liquids are excluded from K expression.

Calculating Equilibrium Concentrations

• Use ICE (Initial, Change, Equilibrium) tables to calculate unknown concentrations:

  • Balanced equation

  • Initial concentrations

  • Changes during reaction

  • Equilibrium concentrations

Effects of Temperature and Pressure Changes

• Increasing Volume decreases pressure, shifting balance toward side with more moles of gas.

• Decreasing Volume increases pressure, shifting balance toward side with fewer moles of gas.

• Temperature Change:

  • Exothermic reactions: Increase in temperature shifts left, decreasing K.

  • Endothermic reactions: Increase in temperature shifts right, increasing K.

Example Problems

• Determining changes in concentrations based on equilibrium shifts when reactants/products are added or removed.

• Assessing direction of reaction using Q vs. K comparisons.

Special Cases and Assumptions

• When initial concentration divided by K is >400, simplifying assumptions may be valid to disregard x in equilibrium expressions.

• Understanding the significance of K values to determine the extent of reactions (larger K: more product present; smaller K: more reactant present).

Summary

• Le Chatelier’s Principle is a fundamental concept in understanding how equilibria respond to changes in concentration, temperature, and pressure, ensuring constant K under unchanged conditions.