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.