Chemical Kinetics and Chemical Equilibrium: Visual Study Guide

Chemical Kinetics

Measuring Reaction Rates

Chemical kinetics is the study of how fast chemical reactions occur and the factors that affect these rates. The rate of a reaction is typically measured as the change in concentration of a reactant or product per unit time.

• Definition: Rate = change in concentration / change in time

• Units: Molarity per second (M/s)

• Experimental Methods: Reaction rates can be measured using various techniques, such as monitoring pressure changes, spectrophotometry, or gas chromatography.

Visualizing Reaction Progress

Reaction progress can be visualized using molecular models or concentration vs. time graphs. The rate of a reaction can be fast or slow, and the concentration of reactants decreases while products increase over time.

• Fast reactions: Rapid decrease in reactant concentration.

• Slow reactions: Gradual decrease in reactant concentration.

Reaction Order and Rate Laws

The rate law expresses the relationship between the rate of a reaction and the concentration of its reactants. The general form is:

• General Rate Law: $\text{Rate} = k[A]^n$

• Zero Order: Rate is independent of [A].

• First Order: Rate is directly proportional to [A].

• Second Order: Rate is proportional to [A]2.

Determining Reaction Order

Reaction order is determined experimentally, often using the method of initial rates. By comparing how the initial rate changes with varying concentrations, the order with respect to each reactant can be deduced.

Chemical Equilibrium

The Concept of Dynamic Equilibrium

In a reversible reaction, dynamic equilibrium is reached when the rate of the forward reaction equals the rate of the reverse reaction. At this point, the concentrations of reactants and products remain constant, though not necessarily equal.

• Dynamic: Both forward and reverse reactions continue to occur.

• Not static: Concentrations are constant, but molecules are still reacting.

The Equilibrium Constant (K)

The equilibrium constant, K, quantifies the ratio of product to reactant concentrations at equilibrium, each raised to the power of their stoichiometric coefficients.

• General Expression: $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

• K > 1: Products favored at equilibrium.

• K < 1: Reactants favored at equilibrium.

Visualizing Equilibrium Mixtures

Equilibrium can be illustrated using particle diagrams, showing the relative amounts of reactants and products at equilibrium for different reactions or conditions.

Le Châtelier’s Principle

Le Châtelier’s Principle states that if a system at equilibrium is disturbed, it will shift in the direction that minimizes the disturbance. Disturbances include changes in concentration, pressure/volume, or temperature.

• Adding reactant: Shifts equilibrium toward products.

• Removing product: Shifts equilibrium toward products.

• Increasing pressure (decreasing volume): Shifts toward side with fewer moles of gas.

• Increasing temperature: Shifts in the direction that absorbs heat (endothermic direction).

Visualizing Shifts in Equilibrium

Changes in equilibrium can be visualized with molecular models showing the effect of adding or removing reactants/products, or changing conditions.

Summary Table: Key Relationships

Change	System Response
Add reactant	Shifts toward products
Remove reactant	Shifts toward reactants
Add product	Shifts toward reactants
Remove product	Shifts toward products
Increase pressure (decrease volume)	Shifts toward fewer moles of gas
Decrease pressure (increase volume)	Shifts toward more moles of gas
Increase temperature (exothermic)	Shifts toward reactants
Increase temperature (endothermic)	Shifts toward products

Key Terms

• Rate Law: Mathematical relationship between reaction rate and reactant concentrations.

• Equilibrium Constant (K): Ratio of product to reactant concentrations at equilibrium.

• Le Châtelier’s Principle: System shifts to minimize disturbances to equilibrium.