Chapter 14: Chemical Kinetics

Overall and Elementary Reactions

Chemical reactions can be described as overall reactions or as a series of elementary steps. Understanding the distinction is crucial for analyzing reaction mechanisms and rates.

• Overall Reaction: The net chemical equation representing the transformation from reactants to products.

• Elementary Reaction: A single step in a reaction mechanism, representing a specific molecular event.

• Molecularity: The number of reactant particles involved in an elementary step (unimolecular, bimolecular, termolecular).

• Reaction Mechanism: The sequence of elementary steps that make up the overall reaction.

• Example: For the reaction , the mechanism may involve two elementary steps.

Reaction Rates and Rate Laws

The rate of a chemical reaction measures how quickly reactants are converted to products. Rate laws express this quantitatively.

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

• Average Rate: Calculated over a finite time interval.

• Instantaneous Rate: The rate at a specific moment, found using calculus (slope of tangent to concentration vs. time curve).

• Empirical Rate Law: Determined experimentally; relates rate to concentrations of reactants.

• General Rate Law:

• Rate Constant (k): Proportionality constant specific to a reaction at a given temperature.

• Reaction Order: The sum of exponents in the rate law; can be zero, first, second, etc.

• Example: For , if rate , the reaction is second order in A.

Determining Rate Laws

Experimental data is used to determine the rate law and reaction order.

• Initial Rates Method: Measure initial rates at different reactant concentrations to deduce the rate law.

• Rate Data in Tables: Tabulated experimental data showing how rate changes with concentration.

Integrated Rate Laws and Half-Life

Integrated rate laws relate concentration to time for different reaction orders.

• First Order:

• Second Order:

• Zero Order:

• Half-Life (): Time for concentration to decrease by half. For first order:

Collision Model and Activation Energy

The collision model explains how molecular collisions lead to reactions, emphasizing the importance of activation energy.

• Activation Energy (): Minimum energy required for a reaction to occur.

• Arrhenius Equation:

• Temperature Dependence: Higher temperature increases rate by providing more energy to surpass .

Catalysis

Catalysts increase reaction rates by providing alternative pathways with lower activation energy, without being consumed.

• Effect on Rate: Increases rate without affecting equilibrium position.

• Types: Homogeneous (same phase as reactants), heterogeneous (different phase).

Chapter 15: Chemical Equilibrium

Equilibrium Concepts and Calculations

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

• Equilibrium Constant (): for

• Homogeneous Equilibria: All species in the same phase.

• Heterogeneous Equilibria: Species in different phases; solids and pure liquids are omitted from expressions.

• Calculating Equilibrium Concentrations: Use ICE tables (Initial, Change, Equilibrium) and solve for unknowns.

Dependence on Stoichiometry and Manipulation of

• Doubling Reaction Coefficients: is squared.

• Reversing Reaction: is inverted.

• Adding Reactions: Multiply the values.

Le Châtelier's Principle

Predicts how a system at equilibrium responds to disturbances (changes in concentration, pressure, volume, or temperature).

• Concentration: Adding/removing reactants or products shifts equilibrium to oppose the change.

• Pressure/Volume: Affects equilibria involving gases; decreasing volume favors side with fewer moles of gas.

• Temperature: Increasing temperature favors endothermic direction; decreasing favors exothermic.

• Catalysts: Do not affect equilibrium position, only the rate at which equilibrium is reached.

Reaction Quotient (Q)

Used to determine the direction a reaction will proceed to reach equilibrium.

• Definition: Same form as but with initial (not equilibrium) concentrations.

• Comparison: If , reaction proceeds forward; if , reaction proceeds in reverse.

Chapter 16: Acid-Base Equilibria

Acid-Base Definitions

• Arrhenius: Acids produce in water; bases produce .

• Brønsted-Lowry: Acids donate protons (); bases accept protons.

• Lewis: Acids accept electron pairs; bases donate electron pairs.

Strong and Weak Acids/Bases

• Strong Acids/Bases: Completely ionize in solution (e.g., , ).

• Weak Acids/Bases: Partially ionize; equilibrium exists between ionized and unionized forms.

• Conjugate Acid-Base Pairs: Differ by one proton.

pH, pOH, and Ionization Constants

• pH:

• pOH:

• Relationship: (at 25°C)

• Acid Dissociation Constant ():

• Base Dissociation Constant ():

• Relationship:

• p and p: ,

Autoionization of Water

• Equation:

• Ion Product: at 25°C

Polyprotic Acids and Bases

• Definition: Acids/bases that can donate/accept more than one proton (e.g., , ).

• Ionization Steps: Each step has its own or value, decreasing with each proton lost/gained.

Acid-Base Properties of Salt Solutions

• Hydrolysis: Salts from weak acids/bases can affect pH of solution.

• Example: forms acidic solution; forms basic solution.

Percent Ionization and Simplifying Assumptions

• Percent Ionization:

• Assumptions: For weak acids/bases, if or is small, (change in concentration) can be neglected in denominator for simplification.

Factors Affecting Acid/Base Strength

• Bond Strength: Weaker bonds to H favor stronger acids.

• Electronegativity: More electronegative atoms stabilize conjugate base, increasing acid strength.

• Resonance: Delocalization of charge stabilizes conjugate base.

Amphoteric and Amphiprotic Substances

• Amphoteric: Can act as acid or base (e.g., ).

• Amphiprotic: Can donate or accept a proton (e.g., ).

Chapter 17: Additional Aspects of Aqueous Equilibria

Common-Ion Effect

The presence of a common ion suppresses the ionization of a weak acid or base.

• Example: Adding to solution decreases $HF$ ionization due to increased .

Buffers and the Henderson-Hasselbalch Equation

• Buffer: Solution that resists changes in pH upon addition of small amounts of acid or base; contains weak acid and its conjugate base or vice versa.

• Henderson-Hasselbalch Equation:

Titrations

• Definition: Gradual addition of one solution to another to determine concentration or analyze reaction progress.

• Equivalence Point: Point at which stoichiometric amounts of acid and base have reacted.

• Indicator: Substance that changes color at (or near) the equivalence point.

Solubility and Precipitation

• Solubility Product (): for

• Precipitation: Occurs when ionic product exceeds .

• Common-Ion Effect: Decreases solubility of a salt in solution containing a common ion.

Complex Ions and Qualitative Analysis

• Complex Ion: Ion formed from a metal ion and one or more ligands (Lewis bases).

• Effect: Formation of complex ions can increase solubility of certain salts.

• Qualitative Analysis: Systematic identification of ions in solution using selective precipitation and complex formation.

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