Equilibrium

Unit 6: Equilibrium

Introduction to Chemical Equilibria

• Chemical equilibria are vital in biological and environmental processes.
• Equilibria involving O2 molecules and hemoglobin are crucial for oxygen transport in the body.
• Carbon monoxide (CO) binds to hemoglobin similarly, causing toxicity.

Dynamics of Liquid and Vapor Equilibrium

• In a closed container, when liquid evaporates, molecules with higher kinetic energy escape into the vapor phase.
• Conversely, vapor molecules may collide with the liquid surface and return to the liquid phase.
• Equilibrium is reached when the rate of evaporation equals the rate of condensation.
  • This is denoted as: $H<em>2O (l) \rightleftharpoons H</em>2O (vap)$
• The system's activity at equilibrium is not static; molecular exchanges occur continuously.
• The equilibrium mixture consists of both reactants and products.

Types of Reactions at Equilibrium

• Chemical reactions can vary in speed (fast or slow), affected by experimental conditions and reactant nature.
• At a certain temperature in a closed vessel, initial changes lead to reactants decreasing and products increasing until equilibrium occurs.
• Dynamic equilibrium is achieved when the concentration of reactants and products no longer changes.

Classification Based on Reaction Extent

• Reactions nearing completion: Negligible reactant concentration remains.
• Reactions with minimal product formation: Majority of reactants remain unchanged at equilibrium.
• Reactions with comparable concentrations of reactants and products: All species are present at significant levels.
• Reaction extent is influenced by concentration and temperature.

Equilibrium in Physical Processes

6.1 Solid-Liquid and Liquid-Vapor Equilibrium

6.1.1 Solid-Liquid Equilibrium

• Ice and water in a thermos at 273K demonstrate equilibrium with no mass change over time.
  • Molecules transfer between phases at equal rates, leading to a stable, dynamic equilibrium.
  • The atmosphere can affect the melting and freezing points depending on pressure.

6.1.2 Liquid-Vapor Equilibrium

• When water is placed in a closed environment with a drying agent, evaporation occurs until equilibrium vapor pressure stabilizes with temperature.
• Different liquids exhibit varying vapor pressures at the same temperature, affecting volatility and boiling points.

6.1.3 Solid-Vapor Equilibrium

• Example with iodine sublimes, creating a dynamic atmosphere of vapor around solid; equilibrium is established without changes in mass.
  • Other examples include camphor and ammonium chloride.

6.1.4 Dissolution Equilibrium

• Saturated solutions maintain a dynamic equilibrium where the rate of dissolution matches crystallization.
• Radioactive sugar can illustrate the dynamic exchange between phases.

6.1.5 General Characteristics of Physical Equilibria

• Equilibrium occurs in closed systems; measurable properties remain constant.
• Physical interactions adjusting to reach equilibrium are dynamic but stable.

Chemical Equilibrium - Dynamic Nature

• Similar to physical processes, chemical reactions can attain an equilibrium state.
• Reactions may proceed in both forward and backward directions, leading to concentration constancy.
  • For example: $A + B \rightleftharpoons C + D$
• Haber’s process illustrates dynamic equilibrium through consistent analysis of ammonia concentration during synthesis.

Importance of Reaction Quotient Q

• Q enables prediction of direction toward equilibrium; if Qc > Kc, the reaction favors reactants; if Qc < Kc, it favors products.
• The equilibrium constant (Kc) relates concentrations of reactants and products at equilibrium.
• Kc can also be expressed through partial pressures (Kp) for gaseous reactions.

Law of Chemical Equilibrium and Equilibrium Constant

Definition of Equilibrium Constant

• For a reversible reaction: $K_c = \frac{[C]^c [D]^d}{[A]^a [B]^b}$
• Kc is derived under constant temperature; changes affect its value.
• The equilibrium constant for reverse reactions is reciprocal: $K'c = \frac{1}{Kc}$
• Conditions changing Kc include concentration variation, temperature, pressure, catalysts, and common ions affecting solubility.

Factors Affecting Equilibria

Le Chatelier’s Principle

• A systematic change in concentration, pressure, temperature, etc., shifts the equilibrium to mitigate the effect, emphasizing a dynamic nature.
• Catalysts speed up reactions without shifting equilibrium position.

Concentration Effects

• Addition/removal of reactants/products changes favoring respective directions to restore equilibrium.
  • Example: In $H<em>2 + I</em>2 \rightleftharpoons 2HI$, adding H2 drives equilibrium right.
• Example of color change with Fe3+ and SCN– demonstrates concentration effect.

Pressure Changes

• Adjusting gas volume alters concentrations and pressures affecting equilibrium leaning toward lesser moles of gas.

Inert Gas Addition

• Inert gases do not affect equilibrium in closed systems; concentration and partial pressures remain largely unchanged.

Temperature Changes

• Temperature increase favors endothermic reactions on a general basis, while exothermic reactions are favored at lower temperatures.
  • Example: Carbon monoxide’s reaction to form dihydrogen in the industry.

Catalyst Impact

• Catalysts expedite reaction rates but do not influence the final equilibrium concentrations of reactants or products.

Ionic Equilibrium in Solution

Electrolytes and Their Classification

• Electrolytes conduct electricity in aqueous solutions as a result of ion dissociation.
• Strong electrolytes are nearly fully ionized; weak electrolytes are not.

Ionization Concepts: Arrhenius, Brönsted-Lowry, Lewis

• Arrhenius acids produce H+ ions; bases produce OH– ions.
• Brönsted-Lowry expands definition recognizing proton donors and acceptors.
• Lewis defines acids as electron pair acceptors, bases as donors, allowing broader classification.

Relationship of Ionization Constants

• Ka and Kb give insight into acid/base strengths; conjugate pairs have inverse relationships.
  • Example: $K<em>a × K</em>b = K_w$
• The pH scale assists in understanding hydrogen ion concentration in solutions.

Buffer Solutions

• Buffers resist pH changes upon dilution or addition of acids/bases.
  • They are often composed of weak acids and their salts.
• The Henderson-Hasselbalch equation offers a method to calculate buffer pH based on pKa and constituent concentrations.

Summary

• Reactions at equilibrium are dynamic. The equilibrium constant reflects concentrations at a fixed temperature, but are influenced by concentration, pressure, temperature, and catalysis.
• Ionic equilibria and buffer systems help manage pH in various contexts, from physiological to industrial.