Equilibrium and Kinetics

Equilibrium Reactions

Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products over time.

• Dynamic Equilibrium: Both forward and reverse reactions continue to occur, but there is no net change in concentrations.

• Equilibrium Constant (K): The ratio of product concentrations to reactant concentrations, each raised to the power of their coefficients in the balanced equation.

• General Expression:

• Example: For the reaction , the equilibrium constant is as shown above.

Homogeneous vs. Heterogeneous Equilibria

Equilibria can be classified based on the physical states of the reactants and products.

• Homogeneous Equilibrium: All reactants and products are in the same phase (usually all gases or all aqueous solutions).

• Heterogeneous Equilibrium: Reactants and/or products are in different phases (e.g., solids, liquids, gases).

• Key Point: Pure solids and liquids are not included in the equilibrium constant expression.

• Example: — Only appears in .

Manipulating Kc

The equilibrium constant () can change depending on how the chemical equation is manipulated.

• Reversing the Equation: Inverts the value of ().

• Multiplying the Equation by n: Raises to the nth power ().

• Adding Equations: The overall is the product of the individual values.

• Example: If for is 4, then for , .

Le Chatelier's Principle

Le Chatelier's Principle predicts how a system at equilibrium responds to disturbances.

• Concentration: Increasing the concentration of a reactant or product shifts equilibrium to oppose the change.

• Pressure/Volume (for gases): Increasing pressure (by decreasing volume) shifts equilibrium toward the side with fewer moles of gas.

• Temperature: Raising temperature favors the endothermic direction; lowering favors exothermic.

• Example: For , increasing shifts equilibrium right (toward ).

ICE Tables

ICE (Initial, Change, Equilibrium) tables are used to organize and solve equilibrium problems.

• Initial: List starting concentrations.

• Change: Show how concentrations change as the system moves toward equilibrium (usually using variables like x).

• Equilibrium: Express final concentrations in terms of initial values and x.

• Example: For , if M, M, and x mol/L of A reacts, then at equilibrium: , .

Gibbs Free Energy and Equilibrium Constants

The relationship between Gibbs free energy and the equilibrium constant determines the spontaneity of a reaction.

• Gibbs Free Energy Change (): Indicates whether a reaction is spontaneous under standard conditions.

• Relationship to :

• Where: R = gas constant, T = temperature in Kelvin, K = equilibrium constant.

• Interpretation: If , is negative (spontaneous); if , is positive (non-spontaneous).

Reaction Rates and Kinetics

Factors that Affect Reaction Rates

The speed of a chemical reaction depends on several factors.

• Concentration: Higher concentrations generally increase reaction rates.

• Temperature: Raising temperature increases kinetic energy and reaction rate.

• Catalysts: Catalysts lower activation energy, increasing rate without being consumed.

• Surface Area: Greater surface area (for solids) increases rate.

• Nature of Reactants: Some substances react faster due to their chemical properties.

• Example: Powdered zinc reacts faster with acid than a zinc strip due to higher surface area.

Reaction Rates

Reaction rate is the change in concentration of a reactant or product per unit time.

• Average Rate:

• Instantaneous Rate: The rate at a specific moment, found by the slope of the concentration vs. time curve.

• Units: Usually mol L-1 s-1.

• Example: If decreases from 1.0 M to 0.8 M in 10 s, average rate = M/s.

Concentrations and Rate Laws

Rate laws express the relationship between the rate of a reaction and the concentrations of reactants.

• General Form:

• Order of Reaction: The exponents m and n indicate the order with respect to each reactant.

• Overall Order: Sum of the exponents (m + n).

• Example: If , the reaction is second order in A, first order in B, third order overall.

Concentration Changes with Time

The integrated rate law describes how reactant concentrations change over time for different reaction orders.

• First Order:

• Second Order:

• Zero Order:

• Half-life: The time required for half the reactant to be consumed; depends on reaction order.

• Example: For a first-order reaction with s-1, s.

Temperature, Rate, and Catalysis

Temperature and catalysts significantly influence reaction rates.

• Arrhenius Equation: Shows the relationship between rate constant and temperature.

• Where: = activation energy, R = gas constant, T = temperature (K), A = frequency factor.

• Effect of Temperature: Increasing T increases k, thus increasing rate.

• Catalysts: Lower , increasing k and rate.

• Example: Enzymes are biological catalysts that speed up reactions in living organisms.

Summary Table: Key Concepts in Equilibrium and Kinetics

Additional info: These notes are based on the provided study guide structure and standard General Chemistry II curriculum. Specific worked examples and detailed video content are not included, as the original file did not provide them.