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EQUILIBIRUM

Dynamic Nature of Equilibrium in Reactions

Equilibrium in both physical and chemical processes is dynamic, characterized by continuous activity where the rates of forward and reverse reactions are equal. This results in constant concentrations of reactants and products despite ongoing reactions.

Equilibrium is established when the rate of evaporation equals the rate of condensation in physical processes.
Chemical reactions can reach equilibrium from either direction, starting with reactants or products.
The dynamic nature of equilibrium is illustrated through examples like the synthesis of ammonia and the use of isotopes.

Characteristics of Solid-Liquid Equilibrium

Solid-liquid equilibrium occurs at a specific temperature and pressure, where both phases coexist without changing mass over time. This equilibrium is dynamic, with continuous molecular exchange between solid and liquid states.

Ice and water at 273K and atmospheric pressure exemplify solid-liquid equilibrium.
The melting point is the temperature at which solid and liquid phases coexist at 1 atm.
The rates of melting and freezing are equal, maintaining constant mass.

Understanding Liquid-Vapour Equilibrium

Liquid-vapour equilibrium is established when the rate of evaporation equals the rate of condensation, resulting in a constant vapour pressure at a given temperature. Different liquids exhibit varying vapour pressures, indicating volatility.

The equilibrium vapour pressure of water is constant at a given temperature.
Boiling point is defined as the temperature at which liquid and vapour phases coexist at 1.013 bar.
Vapour pressure increases with temperature, and different liquids have distinct vapour pressures.

Solid-Vapour Equilibrium Dynamics

Solid-vapour equilibrium occurs when solids sublime into vapour, reaching a state where the rate of sublimation equals the rate of condensation. This dynamic equilibrium is characterized by constant intensity of vapour color over time.

Solid iodine in a closed vessel demonstrates solid-vapour equilibrium.
Other examples include camphor and ammonium chloride.
The intensity of vapour color stabilizes, indicating equilibrium.

Equilibrium in Dissolution Processes

Equilibrium in dissolution involves a dynamic balance between solute in solid form and solute dissolved in solution. The solubility of solids and gases in liquids is temperature-dependent and governed by specific laws.

Saturated solutions exist when no more solute can dissolve at a given temperature.
Henry’s law states that gas solubility in a liquid is proportional to the gas pressure above the liquid.
The concentration of dissolved gases decreases with increasing temperature.

General Characteristics of Physical Equilibria

Physical equilibria share common characteristics, including the necessity of a closed system and the constancy of measurable properties. These characteristics highlight the dynamic yet stable nature of equilibrium.

Equilibrium is only possible in a closed system at a specific temperature.
Opposing processes occur at the same rate, maintaining a stable condition.
Measurable properties remain constant, indicating the extent of the physical process before reaching equilibrium.

Equilibrium Mixtures and Their Composition

Equilibrium mixtures consist of reactants and products in a balanced state, where their concentrations remain constant. Understanding the relationships between these concentrations is crucial for industrial chemical synthesis.

An equilibrium mixture is formed when the rates of the forward and reverse reactions are equal.
The equilibrium constant (Kc) relates the concentrations of reactants and products in a reversible reaction.
The law of mass action states that Kc is determined by the ratio of the concentrations of products to reactants, raised to their stoichiometric coefficients.
For the reaction H2(g) + I2(g) ⇌ 2HI(g), Kc = [HI]^2 / ([H2][I2]).
The equilibrium constant is temperature-dependent and varies with changes in reaction conditions.

Determining Equilibrium Concentrations

Equilibrium concentrations can be calculated from initial concentrations using stoichiometry and the equilibrium constant. This process is essential for predicting the outcomes of chemical reactions.

Initial concentrations are used to set up an ICE (Initial, Change, Equilibrium) table for the reaction.
The change in concentration (x) is defined based on the stoichiometry of the reaction.
The equilibrium constant expression is then used to solve for x, allowing for the calculation of equilibrium concentrations.
Quadratic equations may arise in calculations, and the physically meaningful solution must be chosen.

Homogeneous and Heterogeneous Equilibria

Equilibria can be classified as homogeneous, where all reactants and products are in the same phase, or heterogeneous, where they are in different phases. The equilibrium constant expressions differ based on the phase of the substances involved.

Homogeneous equilibria involve all species in the same phase, such as gases or solutions.
Heterogeneous equilibria include solids or liquids, where their concentrations do not appear in the equilibrium constant expression.
For gaseous reactions, Kp can be used, which is based on partial pressures, and is related to Kc by the equation Kp = Kc(RT)Δn.
The equilibrium constant for a reaction is independent of the amounts of pure solids or liquids present.

Factors Affecting Equilibrium Conditions

Various factors can influence the position of equilibrium, including concentration, pressure, temperature, and the presence of catalysts. Le Chatelier's principle helps predict how a system will respond to changes.

Changes in concentration shift the equilibrium to consume the added substance or replenish the removed one.
Pressure changes affect gaseous reactions differently based on the total number of moles of reactants and products.
Adding an inert gas at constant volume does not affect the equilibrium position.
Temperature changes alter the equilibrium constant; exothermic reactions decrease Kc with increased temperature, while endothermic reactions increase Kc.

Relationship Between Gibbs Energy and Equilibrium Constant

The equilibrium constant is directly related to the Gibbs free energy change of a reaction, providing insights into the spontaneity of the reaction.

A negative ΔG indicates a spontaneous reaction, while a positive ΔG suggests non-spontaneity.
At equilibrium, ΔG = 0, leading to the relationship ΔG° = -RT ln K.
The equation K = e^(-ΔG°/RT) shows that a lower ΔG° corresponds to a higher equilibrium constant, indicating a greater concentration of products at equilibrium.

Applications of Equilibrium Constants

Equilibrium constants are used to predict the extent and direction of chemical reactions, as well as to calculate equilibrium concentrations under various conditions.

A high Kc (> 10^3) indicates that products predominate, while a low Kc (< 10 xss=removed>
The reaction quotient (Qc) can be compared to Kc to determine the direction of the reaction.
Initial concentrations can be used to calculate equilibrium concentrations through systematic approaches involving ICE tables and equilibrium expressions.

Occurrence and Properties of Acids, Bases, and Salts

Acids, bases, and salts are prevalent in nature and play crucial roles in various biological and chemical processes. They exhibit distinct properties such as taste, pH changes, and reactions with indicators like litmus paper.

Hydrochloric acid (HCl) is secreted in gastric juice, essential for digestion, at a rate of 1.2-1.5 L/day.
Acetic acid is the main component of vinegar, while citric and ascorbic acids are found in citrus fruits.
Acids taste sour, turn blue litmus red, and react with metals to release hydrogen gas.
Bases taste bitter, feel soapy, and turn red litmus blue; washing soda is a common example.
Salts, such as sodium chloride, are formed from the reaction of acids and bases.

Arrhenius Concept of Acids and Bases

The Arrhenius theory defines acids and bases based on their behavior in water, focusing on the production of hydrogen and hydroxyl ions. This concept, while foundational, has limitations in its applicability.

Acids dissociate in water to produce hydrogen ions (H+), while bases produce hydroxyl ions (OH–).
The ionization of acids like HX can be represented as HX → H+ + X–.
A bare proton (H+) cannot exist freely in solution and forms hydronium ions (H3O+).
The theory does not account for non-aqueous solutions or bases like ammonia that lack hydroxyl groups.

Brönsted-Lowry Acids and Bases

The Brönsted-Lowry theory expands the definition of acids and bases to include proton donors and acceptors, respectively. This theory provides a more comprehensive understanding of acid-base reactions.

Acids are defined as proton donors, while bases are proton acceptors.
Water can act as both an acid and a base, depending on the reaction context.
Conjugate acid-base pairs differ by one proton; for example, HCl and Cl– are a conjugate pair.
Strong acids have weak conjugate bases, and vice versa.

Lewis Acids and Bases

The Lewis theory further broadens the definitions of acids and bases by focusing on electron pair donation and acceptance. This perspective includes species that do not contain protons.

Lewis acids accept electron pairs, while Lewis bases donate electron pairs.
Examples include BF3 acting as a Lewis acid by accepting a lone pair from NH3, a Lewis base.
This theory encompasses a wider range of chemical reactions compared to previous definitions.

Ionization of Acids and Bases

The ionization of acids and base

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