Chemical Kinetics

Factors Influencing Reaction Rates

Chemical kinetics studies the speed at which chemical reactions occur and the factors that affect these rates.

• Concentration: Higher concentration increases collision frequency, raising reaction rate.

• Temperature: Higher temperature increases kinetic energy, leading to more effective collisions.

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

• Surface Area: For heterogeneous reactions, greater surface area increases rate.

• Nature of Reactants: Ionic reactions are generally faster than covalent reactions.

Expressing Reaction Rate

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

• Rate Expression: for a reaction .

• Average Rate:

• Instantaneous Rate: Determined from the slope of concentration vs. time graph at a specific point.

Rate Constant and Rate Law

The rate law relates the rate of reaction to the concentration of reactants.

• Rate Law:

• Rate Constant (k): A proportionality constant; units depend on reaction order.

• Determining Rate Law: Use experimental data to find how rate changes with concentration.

Integrated Rate Laws

Integrated rate laws relate concentrations of reactants to time for different reaction orders.

• First Order:

• Second Order:

• Zero Order:

• Applications: Calculate concentration at any time, time for a fraction to react, or time to reach a certain concentration.

Half-Life

Half-life is the time required for the concentration of a reactant to decrease by half.

• First Order Half-Life:

• Relationship: For first order, half-life is independent of initial concentration.

Activation Energy and Arrhenius Equation

Activation energy is the minimum energy required for a reaction to occur.

• Arrhenius Equation:

• Temperature Effect: Higher temperature increases rate constant.

• Determining : Use rate constants at different temperatures.

Collision Model and Reaction Mechanisms

The collision model explains how molecules must collide with proper orientation and sufficient energy to react.

• Elementary Step: A single event in a reaction mechanism.

• Rate-Determining Step: The slowest step controls overall rate.

• Intermediate: Species formed and consumed during the mechanism.

• Deriving Rate Law: Use mechanism and identify rate-determining step.

Catalysts

Catalysts increase reaction rate by providing an alternative pathway with lower activation energy.

• Homogeneous Catalyst: Same phase as reactants.

• Heterogeneous Catalyst: Different phase from reactants.

• Recognition: Catalyst appears in mechanism but not in overall equation.

Chemical Equilibrium

Equilibrium Expressions

Chemical equilibrium occurs when the rates of forward and reverse reactions are equal.

• Equilibrium Constant (): for

• Heterogeneous Equilibrium: Pure solids and liquids are omitted from the expression.

Evaluating Equilibrium Constants

• Using Concentrations: Substitute equilibrium concentrations into .

• Using Partial Pressures:

• Interconverting and : , where is the change in moles of gas.

Magnitude of Equilibrium Constant

• Large : Product-favored; more products at equilibrium.

• Small : Reactant-favored; more reactants at equilibrium.

Reaction Quotient () and Predicting Direction

• Reaction Quotient: Same form as , but uses initial concentrations.

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

Le Chatelier’s Principle

Le Chatelier’s Principle predicts how equilibrium shifts in response to changes.

• Concentration: Adding reactant shifts equilibrium toward products.

• Temperature: Endothermic: increase shifts toward products; exothermic: increase shifts toward reactants.

• Pressure/Volume (Gases): Increase pressure (decrease volume) shifts toward fewer moles of gas.

Catalysts and Equilibrium

• Effect: Catalysts do not change equilibrium position; they speed up attainment of equilibrium.

Acids and Bases

Properties and Ions

Acids and bases are characterized by their ability to donate or accept protons.

• Acids: Sour taste, react with metals, turn litmus red; produce ions.

• Bases: Bitter taste, slippery feel, turn litmus blue; produce ions.

Bronsted-Lowry Theory

The Bronsted-Lowry theory defines acids as proton donors and bases as proton acceptors.

• Acid: Donates

• Base: Accepts

• Conjugate Acid-Base Pair: Two species differing by one proton.

• Example:

Autoionization of Water

Water can ionize to form and ions.

• Ion-Product Constant: at 25°C

pH and Calculations

pH is a measure of acidity or basicity of a solution.

• Definition:

• pOH:

• Relationship:

• Reverse Calculation:

Strong and Weak Acids/Bases

• Strong Acids: Completely ionize in water (e.g., HCl, HNO3, H2SO4).

• Strong Bases: Completely ionize in water (e.g., NaOH, KOH).

• Weak Acids/Bases: Partially ionize; use or for calculations.

Calculating pH for Weak Acids and Bases

• Weak Acid:

• Weak Base:

• Percent Ionization:

Polyprotic Acids

Polyprotic acids can donate more than one proton; each dissociation has its own .

• Example: has two dissociation steps.

Acid/Base Strength and Conjugates

• Relationship: Strong acid has weak conjugate base; weak acid has stronger conjugate base.

• Calculation:

Salt Solutions

Salts can produce acidic, basic, or neutral solutions depending on their ions.

• Acidic Salt: Contains conjugate acid of a weak base.

• Basic Salt: Contains conjugate base of a weak acid.

• Neutral Salt: Contains ions from strong acid and strong base.

Acid Strength, Polarity, and Bond Strength

• Polarity: More polar H--X bond increases acid strength.

• Bond Strength: Weaker H--X bond increases acid strength.

Lewis Acid-Base Theory

The Lewis theory defines acids as electron pair acceptors and bases as electron pair donors.

• Lewis Acid: Accepts an electron pair.

• Lewis Base: Donates an electron pair.

• Example: (acid) + (base)