Chapter 14: Chemical Kinetics

Introduction to Chemical Kinetics

Chemical kinetics is the study of the rates at which chemical reactions occur and the factors that affect these rates. Understanding kinetics allows chemists to control reactions and optimize conditions for desired outcomes.

• Reaction Rate: The change in concentration of a reactant or product per unit time.

• Average Rate: Calculated over a time interval; Instantaneous Rate: The rate at a specific moment, found as the slope of the concentration vs. time curve at that point.

Rate Laws

• Rate Law: An equation that relates the reaction rate to the concentrations of reactants, generally in the form:

• Order of Reaction: The sum of the exponents in the rate law; can be zero, first, or second order with respect to a reactant.

• Zero Order:

• First Order:

• Second Order: or

Concentration vs. Time Profiles

• Zero Order:

• First Order:

• Second Order:

Half-Life

• Half-life (t1/2): The time required for the concentration of a reactant to decrease by half.

• First Order:

• Zero Order:

• Second Order:

Activation Energy and Reaction Mechanisms

• Activation Energy (Ea): The minimum energy required for a reaction to occur.

• Arrhenius Equation:

• Transition State/Activated Complex: A high-energy, unstable arrangement of atoms at the peak of the activation energy barrier.

• Mechanism: The sequence of elementary steps by which a reaction occurs.

• Catalyst: A substance that increases the reaction rate by lowering the activation energy, without being consumed.

Example: The decomposition of hydrogen peroxide is catalyzed by iodide ions, which provide an alternative pathway with lower activation energy.

Chapter 15: Chemical Equilibrium

Introduction to Equilibrium

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

• Difference from Kinetics: Kinetics deals with how fast reactions proceed; equilibrium describes the final state where no net change occurs.

Equilibrium Constants

• Equilibrium Constant (K): Expresses the ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their coefficients.

• General Form: for

• Kc: Based on concentrations (mol/L); Kp: Based on partial pressures (atm).

• Kp and Kc Relationship: , where is the change in moles of gas.

• Interpretation of K:

  • K >> 1: Products favored at equilibrium

  • K ≈ 1: Significant amounts of both reactants and products

  • K << 1: Reactants favored at equilibrium

• Manipulation of K: Reversing a reaction inverts K; multiplying coefficients raises K to that power.

ICE Tables

• ICE Table: Used to calculate equilibrium concentrations. Stands for Initial, Change, and Equilibrium.

• Finding K: Use initial and equilibrium concentrations to solve for K.

• Finding Equilibrium Concentrations: Use initial concentrations and K to solve for unknowns.

Reaction Quotient (Q) and Predicting Direction

• Reaction Quotient (Q): Calculated like K, but with initial (not equilibrium) concentrations.

• Comparing Q and K:

  • Q < K: Reaction proceeds forward (toward products)

  • Q > K: Reaction proceeds in reverse (toward reactants)

  • Q = K: System is at equilibrium

Le Châtelier’s Principle

• If a system at equilibrium is disturbed, it will shift to minimize the disturbance.

• Disturbances and Responses:

  • Adding Reactants: Shifts equilibrium toward products

  • Removing Reactants/Products: Shifts to replace what was removed

  • Changing Pressure (gases): Increase favors side with fewer moles of gas

  • Changing Volume: Decreasing volume increases pressure, favoring side with fewer moles of gas

  • Changing Temperature: For endothermic reactions, increasing T favors products; for exothermic, favors reactants

  • Adding Catalyst: No effect on position of equilibrium; only increases rate of reaching equilibrium

Example: Increasing the pressure of the reaction shifts equilibrium toward ammonia, which has fewer moles of gas.

Chapter 16: Acid-Base Equilibria

pH, pOH, and Ionization Constants

Acid-base equilibria involve the transfer of protons (H+) between species in aqueous solution. The strength of acids and bases is quantified using equilibrium constants and pH scales.

• pH:

• pOH:

• Relationship: (at 25°C)

• Ion Product of Water (Kw): at 25°C

• Acid Dissociation Constant (Ka):

• Base Dissociation Constant (Kb):

Strong and Weak Acids/Bases

• Strong Acids/Bases: Completely ionize in water; pH or pOH can be calculated directly from concentration.

• Weak Acids/Bases: Partially ionize; require equilibrium calculations (often using ICE tables).

• 5% Approximation: If the amount ionized is less than 5% of the initial concentration, the change (x) can be neglected in the denominator for simplification.

Acidity Strength, Bond Strength, and Electronegativity

• Acid Strength: Stronger acids have weaker H–A bonds and/or more electronegative atoms stabilizing the conjugate base.

• Base Strength: Stronger bases have a greater tendency to accept protons.

Polyprotic Acids and Bases

• Polyprotic Acids: Acids that can donate more than one proton (e.g., H2SO4).

• Successive Ionization Constants: , , etc.;

• Polyprotic Bases: Similar concept with , , etc.

ICE Tables in Acid-Base Equilibria

• Used to determine equilibrium concentrations of all species in weak acid/base solutions.

• Set up with initial concentrations, changes (usually -x for reactants, +x for products), and equilibrium values.

P Function

• P Function: General form (e.g., pH, pOH, pKa).

Example: Calculate the pH of a 0.10 M acetic acid solution () using an ICE table and the 5% approximation.

Additional info: The study guide references the use of ICE tables, Le Châtelier’s Principle, and the relationship between acid/base strength and molecular structure, which are all foundational concepts in general chemistry. The above notes expand on these points for clarity and completeness.