equilibrium

Page 1: Introduction to Equilibrium

Importance of Chemical Equilibria

  • Chemical equilibria are crucial in numerous biological and environmental processes

    • Example: O2 molecules and hemoglobin play a key role in the transport and delivery of oxygen.

    • CO molecules interactions with hemoglobin lead to toxicity.

Physical Equilibria

  • Evaporation and Vapour Pressure

    • In a closed container, molecules with higher kinetic energy escape into the vapor phase while others condense back.

    • Equilibrium state is when the number of evaporating molecules equals those condensing back into the liquid, e.g.,H2O (l) ⇌ H2O (vap)

    • Not static; dynamic activities occur at boundaries between phases.

Types of Equilibrium

  • Equilibrium mixtures can arise from physical processes and chemical reactions, classified based on:

    • Rate and extent of reaction.

Stages of Dynamic Equilibrium

  • Dynamic equilibrium involves:

    1. Forward reaction rate = Reverse reaction rate

    2. Reactant concentrations decrease while product concentrations increase until equilibrium is reached.

Goals of Study

In this unit, you will learn to:

  1. Identify dynamic nature of equilibrium.

  2. State the law of equilibrium.

  3. Explain characteristics of equilibria.

  4. Write expressions for equilibrium constants.

  5. Understand the relationship between Kp and Kc.

  6. Describe classification of acids and bases (Arrhenius, Bronsted-Lowry, Lewis).

  7. Differentiate between weak and strong acids/bases.

  8. Explain factors affecting ionization.

  9. Discuss pH scale and ionization of water.

  10. Appreciate buffer solutions.

  11. Calculate solubility product constant.

Page 2: Classification of Reactions

Types of Reactions at Equilibrium

  1. Reactions nearly completing

    • Only negligible concentrations of reactants remain.

  2. Reactions with small product amounts

    • Most reactants stay at equilibrium.

  3. Reactions with comparable concentrations

    • At equilibrium, concentrations of products and reactants are similar.

Factors Influencing Equilibrium

  • Equilibrium extent varies with experimental conditions such as concentration and temperature. Understanding these conditions is vital for optimization in industrial and laboratory settings.

Physical Equilibrium Examples

Solid-Liquid Equilibrium

  • Example: Ice and water in an insulated thermos at 273K and 1 atm.

    • No mass change; dynamic equilibrium evident at the phase boundary.

    • Temperature and pressure dictate the equilibrium state.

Liquid-Vapor Equilibrium

  • Example: Water in a sealed box with a drying agent.

    • Observed as changes in mercury level in a manometer until no net evaporation occurs.

    • At equilibrium, the vapor pressure remains constant and varies with temperature.

Page 3: Solid-Vapor Equilibrium to Dissolution

Solid-Vapor Equilibrium

  • Example: Iodine solid subliming to vapor.

    • Eventually, concentration of vapor becomes constant.

    • Dynamic equilibrium results in:I2(s) ⇌ I2(vap)

Dissolution of Solids and Gases in Liquids

Solids in Liquids

  • Saturated solutions: Defined as the maximum solute at a specific temperature.

    • Example: Sugar, where rate of dissolution equals crystallization.

    • Radioactive sugar shows the dynamic exchange during equilibrium.

Gases in Liquids

  • Example: Carbon dioxide dissolution in water.

    • Governed by Henry’s law: mass of gas dissolved is proportional to partial pressure.

    • Opening soda water results in gas effervescence, demonstrating equilibrium.

Page 4: Characteristics of Physical Equilibrium

General Features

  1. Closed systems: Equilibrium exists only in contained environments.

  2. Dynamic conditions: Opposing processes occur simultaneously yet balance.

  3. Consistent properties: All measurable aspects remain constant.

  4. Equilibrium characterizations: Defined by established parameter values at given temperatures.

Page 5: Concept of Dynamic Chemical Equilibrium

Features of Chemical Equilibria

  • Similar to physical equilibria, chemical reactions achieve a state where concentrations remain constant.

  • Dynamic nature: Involves simultaneous forward and reverse reaction rates.

  • Example Reaction: A + B ⇌ C + D

    • Forward constructs products, while reverse regenerates reactants until equilibrium ensues, shown as:

  • This state can be approached from either reactants or products starting point.

Page 6: Demonstrating Dynamic Equilibrium

Example Activity:

  • Using colored water in two measuring cylinders, demonstrate constant levels indicating equilibrium despite continued movement.

    • Highlight differences in tube diameters affecting final levels.

Page 7: Law of Chemical Equilibrium

Equilibrium Constant (Kc)

  • Relates the concentration products and reactants at equilibrium using the formula:

Kc = [C]c[D]d / [A]a[B]b

  • Example: H2 + I2 ⇌ 2HI => Kc = [HI]^2 / [H2][I2]

Page 8: Factors Affecting Equilibrium

Equilibrium States

  • Determines impact of various factors on the equilibrium, such as temperature and concentration

  • Example Experiments: Explore result of varying conditions in reversible reactions to observe equilibrium shifts.

Page 9: Solubility Product Constant

Saturated Solutions

  • Solubility equilibrium Ksp defined with respect to concentration of ions at saturation point.

Ksp = [A]^m[B]^n for a general salt

\text{Example: BaSO}_4(s) ⇌ Ba^{2+}(aq) + SO4^{2-}(aq)

Page 10: Classifications in Acids and Bases

Arrhenius, Bronsted-Lowry, and Lewis Definitions

  • Identification and classification according to definitions provided by scientific foundations such as Arrhenius for acids and bases, Bronsted-Lowry for proton transfer concepts, and Lewis for electron transfer dynamics.