Chemical Equilibrium

Introduction

Chemical equilibrium is a fundamental concept in general chemistry, describing the state in which the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products. This chapter covers the principles, calculations, and applications of chemical equilibrium, including the equilibrium constant, Le Châtelier’s Principle, and the effects of changing reaction conditions.

Table of Contents

• The Concept of Equilibrium

• The Equilibrium Constant

• Equilibrium Constants in Terms of Pressure

• Units of Equilibrium Constants

• Working with Equilibrium Constants

• Manipulating Equilibrium Constants

• Heterogeneous Equilibria

• Calculating Equilibrium Constants

• Applications of Equilibrium Constants

• Predicting the Direction of Reaction

• Calculating Equilibrium Concentrations

• Le Châtelier’s Principle

The Concept of Equilibrium

Dynamic Equilibrium

At dynamic equilibrium, the rate of the forward reaction equals the rate of the reverse reaction. Although both reactions continue to occur, the concentrations of reactants and products remain constant over time.

• Dynamic means that the process is ongoing, not static.

• Equilibrium is represented by a double arrow () in chemical equations.

• Example:

Key Points:

• Both forward and reverse reactions occur simultaneously.

• At equilibrium, the rates of the two reactions are equal.

• The concentrations of all species remain constant (but not necessarily equal).

The Equilibrium Constant

Definition and Expression

The equilibrium constant () quantifies the ratio of product and reactant concentrations at equilibrium for a given reaction at a specific temperature.

• For a general reaction:

The equilibrium constant expression is:

• Square brackets denote molar concentrations (mol/L).

• is used for concentrations; is used for partial pressures (for gases).

Equilibrium Constants in Terms of Pressure

For gaseous reactions, the equilibrium constant can be expressed in terms of partial pressures:

• = partial pressure of species X (in atm or bar).

• Relationship between and : where (change in moles of gas), is the gas constant, and is temperature in Kelvin.

Units of Equilibrium Constants

• The units of and depend on the reaction stoichiometry.

• If , is dimensionless.

• Otherwise, units are derived from the exponents in the equilibrium expression.

Working with Equilibrium Constants

Magnitude and Meaning of

• If , products predominate at equilibrium (product-favored reaction).

• If , reactants predominate at equilibrium (reactant-favored reaction).

• If , significant amounts of both reactants and products are present.

Manipulating Equilibrium Constants

• Reversing a reaction:

• Multiplying a reaction by :

• Adding reactions:

Example Table: Manipulation of Equilibrium Constants

Heterogeneous Equilibria

Definition and Treatment

Heterogeneous equilibria involve reactants and products in different phases (e.g., solids, liquids, gases).

• The concentrations of pure solids and liquids are constant and omitted from the equilibrium expression.

• Only include gases and aqueous species in expressions.

Example:

Equilibrium expression:

Calculating Equilibrium Constants

General Steps

1. Write the balanced chemical equation.

2. Tabulate initial concentrations, changes, and equilibrium concentrations (ICE table).

3. Use stoichiometry to relate changes in concentrations.

4. Substitute equilibrium concentrations into the expression and solve for unknowns.

Example ICE Table

Applications of Equilibrium Constants

• Predict the direction a reaction will proceed to reach equilibrium.

• Calculate the concentrations of reactants and products at equilibrium.

• Assess the extent of a reaction.

Reaction Quotient ()

The reaction quotient () is calculated using the same expression as , but with current (not necessarily equilibrium) concentrations.

• If , the system is at equilibrium.

• If , the reaction proceeds forward (toward products).

• If , the reaction proceeds in reverse (toward reactants).

Le Châtelier’s Principle

Statement and Applications

Le Châtelier’s Principle states: If a system at equilibrium is disturbed by a change in concentration, pressure, or temperature, the system will shift its equilibrium position to counteract the disturbance.

Effects of Changing Conditions

• Concentration: Adding reactant or product shifts equilibrium to consume the added component.

• Pressure/Volume (for gases): Increasing pressure (decreasing volume) shifts equilibrium toward the side with fewer moles of gas; decreasing pressure (increasing volume) shifts toward more moles of gas.

• Temperature:

  • For endothermic reactions (), increasing temperature favors products.

  • For exothermic reactions (), increasing temperature favors reactants.

• Catalysts: Catalysts increase the rate at which equilibrium is achieved but do not affect the equilibrium position or composition.

Summary Table: Le Châtelier’s Principle

Key Equations

• General equilibrium constant:

• For gases:

• Relationship between and :

• Reaction quotient: (using current concentrations)

Example Application: In the Haber process (), removing ammonia as it forms shifts the equilibrium to the right, increasing yield.

Additional info: These notes are based on the provided lecture slides and standard general chemistry textbooks. All equations and tables have been reconstructed for clarity and completeness.