Chemical Equilibrium and Solubility: AP Chemistry Unit 7 Study Guide

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Equilibrium Fundamentals

Introduction to Chemical Equilibrium

Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products over time. This concept is central to understanding reversible reactions in chemistry.

Reversible Processes: Many chemical reactions can proceed in both directions, such as evaporation/condensation, dissolution/precipitation, acid-base, and redox reactions.
Equilibrium Condition: At equilibrium, the concentrations of all species remain constant, though reactions continue to occur at equal rates in both directions.
Graphical Representation: Equilibrium is visually indicated when concentration curves level off.
Example: The reaction between H2 and N2 to form NH3 demonstrates equilibrium when the concentrations of all species stabilize.

Concentration vs. time graph for H2, N2, and NH3

Equilibrium Expressions and Constants

The equilibrium constant (K) quantifies the ratio of product to reactant concentrations at equilibrium. It is specific to each reaction and depends on temperature.

Kc: Based on concentrations (mol/L).
Kp: Based on partial pressures (atm).
Reaction Quotient (Q): Used to determine the direction a reaction will shift to reach equilibrium.
Exclusion of Solids and Liquids: Only gases and aqueous species are included in equilibrium expressions.
Example: For H2(g) + I2(s) ⇄ 2HI(g),

Using RICE Tables

RICE tables (Reaction, Initial, Change, Equilibrium) are a systematic way to solve equilibrium problems, especially when initial and final values are given.

Setup: List initial concentrations, changes (using x), and equilibrium values.
Example: For COCl2(g) ⇄ CO(g) + Cl2(g), if initial pressure is 1.0 atm and equilibrium total pressure is 1.2 atm, use the change in pressure to solve for equilibrium concentrations.

Comparing K Values

The magnitude of K indicates whether products or reactants are favored at equilibrium.

K > 1: Products are favored.
K < 1: Reactants are favored.
Very large K: Reaction proceeds nearly to completion.
Very small K: Reaction barely proceeds.

Manipulating Equilibrium

Relationship Between K and the Chemical Equation

Changes to the chemical equation require adjustments to the equilibrium constant.

Reversing the equation: Invert K.
Multiplying coefficients: Raise K to the power of the factor.
Adding reactions: Multiply the K values of the individual reactions.

Reversing equilibrium equation and inverting K Multiplying coefficients and adjusting K Adding reactions and multiplying K

Solving for Equilibrium Concentrations

Equilibrium concentrations can be calculated using the RICE table and the equilibrium constant expression.

Quadratic Formula: Sometimes required if the equilibrium expression is quadratic.
x-is-small Approximation: Used when x is much smaller than initial concentrations (error < 5%).
Example: For N2(g) + O2(g) ⇄ 2NO(g), , use RICE table and solve for x.

RICE table for N2, O2, and NO equilibrium

Le Chatelier's Principle and Shifts in Equilibrium

Effect of Concentration Changes

Adding or removing reactants/products causes the system to shift to reestablish equilibrium.

Q < K: Reaction shifts right (toward products).
Q > K: Reaction shifts left (toward reactants).
Example: Adding NO2 to N2O4 ⇄ 2NO2 causes Q > K, so the reaction shifts left.

NO2 and N2O4 concentration vs. time graph

Effect of Volume and Pressure

Changing the volume (or pressure) of a gaseous system shifts equilibrium toward the side with fewer or more moles of gas.

Decrease volume (increase pressure): Shifts toward fewer moles of gas.
Increase volume (decrease pressure): Shifts toward more moles of gas.
Example: N2(g) + 3H2(g) ⇄ 2NH3(g) shifts right when volume decreases.

Particle diagram showing volume change and equilibrium shift

Effect of Temperature

Temperature changes affect the equilibrium constant (K), not Q. Heat can be considered a reactant (endothermic) or product (exothermic).

Exothermic: Increase temperature shifts left; decrease shifts right.
Endothermic: Increase temperature shifts right; decrease shifts left.
Van 't Hoff Equation: Relates enthalpy change to the equilibrium constant.

Exothermic and endothermic equilibrium shifts Van 't Hoff equation

Solubility Equilibria

Solubility-Product Constant (Ksp)

Ksp describes the equilibrium between a solid salt and its ions in solution. It is used to predict solubility and calculate molar solubility.

Ksp Expression: For CaF2(s) ⇄ Ca2+(aq) + 2F-(aq),
Magnitude: Higher Ksp means more soluble; lower Ksp means less soluble.
Solubility Table: Qualitative and quantitative data for common salts.

Solubility product constants table

Molar Solubility Calculations

Molar solubility is the number of moles of solute that can dissolve per liter before saturation. It can be calculated from Ksp using RICE tables.

Example: For PbCl2(s) ⇄ Pb2+(aq) + 2Cl-(aq), , solve for x.

RICE table for PbCl2 solubility

Common Ion Effect

The solubility of a salt decreases when a solution already contains one of its ions. This is explained by Le Chatelier's Principle.

Example: CaF2(s) in a solution with NaF (common ion F-) will be less soluble.
Calculation: Adjust the Ksp expression to include the initial concentration of the common ion.

RICE table for CaF2 solubility with common ion

Effect of pH on Solubility

Solubility is affected by pH when one of the ions is a weak acid or base. Increasing acidity (lower pH) increases solubility of salts with basic anions.

Example: Mg(OH)2(s) ⇄ Mg2+(aq) + 2OH-(aq) becomes more soluble as pH decreases.

Multiple choice: effect of pH on Mg(OH)2 solubility

Thermodynamics and Equilibrium

Gibbs Free Energy (ΔG°)

Gibbs free energy change determines whether a reaction is spontaneous and relates to equilibrium. It combines enthalpy and entropy:

Equation:
Spontaneity: Negative ΔG means spontaneous; positive ΔG means non-spontaneous.
Entropy (S): A measure of disorder; higher entropy means more disorder.

Free Energy and Dissolution

The free energy change for dissolution reflects the breaking of intermolecular forces, reorganization of solvent, and interaction of dissolved species. Predicting the total ΔG can be complex due to competing factors.

Practice and Application

Particle Diagrams and Solubility

Particle diagrams help visualize the dissolution process and the relative solubility of compounds.

Example: Comparing FeCO3 and MnCO3 solubility using particle diagrams and Ksp values.

Particle diagram for FeCO3 and MnCO3 dissolution Table comparing FeCO3 and MnCO3 Ksp, ΔH, and ΔS

Entropy and Dissolution

When ionic compounds dissolve, entropy usually increases as ions disperse in solution.

Example: Dissolution of NH4NO3 in water increases entropy.

Particle diagram for NH4NO3 dissolution

Summary Table: Solubility Guidelines

Solubility rules help predict whether a compound will dissolve in water.

Solubility guidelines for common ionic compounds

Summary Table: Solubility Product Constants

Ksp values for common salts provide quantitative data for solubility calculations.

Solubility product constants table

Summary Table: FeCO3 and MnCO3 Dissolution

Comparison of Ksp, ΔH, and ΔS for FeCO3 and MnCO3 dissolution.

Table comparing FeCO3 and MnCO3 Ksp, ΔH, and ΔS Additional info: This study guide covers all major aspects of chemical equilibrium and solubility as outlined in AP Chemistry Unit 7, including graphical, tabular, and particle diagram representations, and provides practice questions for application.