Chemical Equilibrium Concepts

General Equilibrium

• Many chemical reactions are reversible and proceed in both the forward and reverse directions.
• Reversible reactions are denoted using a double-headed arrow, for example, N₂O₄(g) ⇌ 2NO₂(g).
• Forward and reverse reactions are constantly occurring.
• Equilibrium is reached when there's no net change in reactants and products.
• Equilibrium is dynamic, meaning forward and reverse reactions continue, but at equal rates.
• Equilibrium can only be reached in closed systems.
• At equilibrium, the forward and reverse reaction rates are equal, but reactant and product amounts are not necessarily equal.
• Equilibrium is constant at a specific temperature.
• Equilibrium can be visually represented as a concentration vs time graph for a reversible reaction, where slopes are zero at equilibrium.

Chemical Equilibrium

• The forward and reverse reactions do not stop when equilibrium is reached; rather, they occur at equal rates.
• Chemical Equilibrium is dynamic.
• At equilibrium, the rates of the forward and reverse processes are the same, but the amounts of reactants and products might not be equal.
• Equilibrium is a state where there is no observable change.

Representing Equilibrium Graphically

• Determining equilibrium graphically involves constructing a concentration vs time graph for a reversible reaction.

Quantitative Look at Equilibrium

• The equilibrium constant (Keq) represents the concentrations of products and reactants at equilibrium at a given temperature.
• The equilibrium expression is the ratio of product concentrations raised to their stoichiometric coefficients to reactant concentrations raised to their stoichiometric coefficients.
• Solids and liquids are excluded from the equilibrium expression.

Effects of Stresses on Equilibrium

• When equilibrium is stressed (by changes in concentration, volume, pressure, temperature), the system shifts to relieve that stress.

• Increasing the concentration of a product shifts the equilibrium toward reactants.

• Increasing the concentration of a reactant shifts the equilibrium toward products.

• Increasing the volume/decreasing the pressure shifts the equilibrium to the side with more gas molecules.

• Decreasing the volume/increasing the pressure shifts the equilibrium to the side with fewer gas molecules.

• Increasing the temperature shifts the equilibrium to the endothermic direction.

• Decreasing the temperature shifts the equilibrium to the exothermic direction.

• The addition of a catalyst does not affect the position of equilibrium.

• The addition of an inert gas at constant volume does not affect the position of equilibrium.

• The addition of an inert gas at constant pressure shifts the equilibrium to the side with more moles of gas.

ICE Tables

• Ice tables are used to solve equilibrium problems by organizing information about initial concentrations, calculated changes, and equilibrium concentrations of reactants and products.

Solubility Equilibrium

• A solution is a homogeneous mixture of one or more solutes dissolved in a solvent.
• Solutes are dissolved substances.
• A solvent is the dissolving substance.
• Solutions can have variable composition.
• Aqueous solutions contain water as a solvent.
• Types of solutes include ionic compounds and polar and non-polar molecules.
• Solubility is the amount of solute that dissolves in a given quantity of solvent at a particular temperature.
• Saturated solutions are those in which no more of a particular solute can be dissolved.
• Unsaturated solutions are those in which more of a solute can be dissolved.

Solubility Equilibrium (Continued)

• Important Factors Affecting Solubility •Molecule size (larger size, lower solubility) •Temperature (higher temperature, increases solubility in solids, gases; more collisions and higher solubility in gases) •Pressure (higher pressure, higher solubility in gases).
• Concepts •Soluble – Solubility > 1 g / 100 mL of solvent •Insoluble – Solubility < 0.1 g / 100 mL of solvent •Slightly soluble – Solubility between 0.1 g and 1 g / 100 mL of solvent
• Like dissolves like

Equilibrium and Le Chatelier's Principle

• Predict shifts and effects on solubility equations (e.g., PbI₂(s) ↔ Pb²⁺(aq) + 2I⁻(aq)).
• Selective precipitation involves manipulating conditions to force the precipitation of one ion while leaving others dissolved, considerations of salt solubility are important.

Arrhenius vs Brønsted-Lowry Acids and Bases

• Arrhenius acids yield H⁺ ions in solution; Arrhenius bases yield OH⁻ ions in solution.
• Brønsted-Lowry acids are proton (H⁺) donors; Brønsted-Lowry bases are proton (H⁺) acceptors.
• Conjugate acids and bases are formed when bases gain a hydrogen ion (acid donates), or acids lose a hydrogen ion(base accepts).
• Conjugate acid-base pairs are formed when an acid and a base are in equilibrium.

Strong vs Weak Acids and Bases

• Strong acids/bases ionize completely in water (e.g., HCl, NaOH).
• Weak acids/bases ionize partially (e.g., CH₃COOH, NH₃).
• The equilibrium constant for weak acids is Ka, and the equilibrium constant for weak bases is Kb.

Autoionization of Water

• Water self-ionizes (to form H₃O⁺ and OH⁻ ions).
• The equilibrium expression is Keq = [H₃O⁺][OH⁻] = 1.0 x 10⁻¹⁴.

Amphoteric Substances

• Substances that can behave as both acids and bases (e.g., HCO3⁻).

Polyprotic Acids

• Acids with more than one ionizable proton (eg. H₂SO₄, H₃PO₄).

Acid-Base Properties of Salt Solutions

• Salts can form neutral, acidic, or basic solutions based on the conjugate acid or base properties of the ion.

Acid-Base Titration

• Indicators are used to determine equivalence point of a titration.
• Equivalence point is reached when chemically equivalent amounts of acid and base have reacted.
• Equivalence point pH depends on the type of reaction.
• Calculations for strong acid/strong base, weak acid/strong base, and strong acid/weak base titrations typically use ICE tables.
• Titration curves are plotted as pH vs. volume of acid/base added.
• The inflection point is where equivalence point occurs.

Finding pH at the Midpoint

• At the midpoint, the concentration of the weak acid is equal to that of its conjugate base.
• pH = pKa