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AP Chemistry

Study notes
1.1.1 Understanding the Mole Concept
1.1.2 Avogadro’s Number and Molar Mass
1.1.3 Mass and Number of Particles
1.2.1 Understanding Mass Spectroscopy
1.2.2 Isotopes Identification through Mass Spectroscopy
1.2.3 Calculating Average Atomic Mass
1.3.1 Introduction to Elemental Composition
1.3.2 Law of Definite Proportions
1.3.3 Empirical Formulas
1.4.1 Understanding Mixtures
1.4.2 Types of Mixtures
1.4.3 Elemental Composition of Mixtures
1.5.1 Composition of the Atom
1.5.2 Coulomb’s Law in Atomic Structure
1.5.3 Electron Shells and Subshells
1.5.4 The Aufbau Principle and Electron Configuration
1.5.5 Ionization Energy and Coulomb’s Law
1.6.1 Introduction to Photoelectron Spectroscopy
1.6.2 Analyzing PES Spectra
1.6.3 Electron Configuration and PES
1.6.4 Electron-Nucleus Interactions
1.7.1 Organization of the Periodic Table
1.7.2 Understanding Periodicity
1.7.3 Trends in Ionization Energy
1.7.4 Trends in Atomic and Ionic Radii
1.7.5 Trends in Electron Affinity and Electronegativity
1.7.6 Predicting Properties in the Absence of Data
1.8.1 Valence Electrons and Chemical Reactivity
1.8.2 Periodicity and Formation of Analogous Compounds
1.8.3 Typical Charges and Ionic Compounds
2.1.1 Electronegativity and Bond Type
2.1.2 Nonpolar Covalent Bonds
2.1.3 Polar Covalent Bonds
2.1.4 Ionic vs. Covalent Bonding
2.1.5 Metallic Bonding
2.2.1 Potential Energy and Equilibrium Bond Length
2.2.2 Bond Length and Bond Order
2.2.3 Coulomb’s Law and Ionic Interactions
2.3.1 Particulate Model of Ionic Solids
2.3.2 Coulomb's Law in Ionic Crystals
2.3.3 Properties of Ionic Solids
2.4.1 Metallic Bonding Model
2.4.2 Structure and Properties of Alloys
2.4.3 Interstitial Alloys
2.4.4 Substitutional Alloys
2.5.1 Principles of Lewis Diagrams
2.5.2 Drawing Lewis Diagrams for Simple Molecules
2.5.3 Lewis Diagrams for Polyatomic Ions
2.5.4 Exceptions to the Octet Rule
2.6.1 Understanding Resonance in Lewis Diagrams
2.6.2 Drawing Resonance Structures
2.6.3 Formal Charge and Its Application
2.6.4 Selecting the Best Lewis Structure
2.6.5 Limitations of Lewis Structures
2.7.1 Basics of VSEPR Theory
2.7.2 Predicting Molecular Geometry with VSEPR
2.7.3 Understanding Bond Hybridization
2.7.4 Application of Hybridization to Molecular Structure
2.7.5 Sigma and Pi Bonds in Molecular Structure
3.1.1 London Dispersion Forces
3.1.2 Polarizability and Molecular Size
3.1.3 Dipole Interactions
3.1.4 Hydrogen Bonding
3.1.5 Noncovalent Interactions in Biomolecules
3.2.1 Intermolecular Forces and Solid Properties
3.2.2 Particulate-Level Structure and Macroscopic Properties
3.2.3 Properties of Ionic Solids
3.2.4 Covalent Network Solids
3.2.5 Molecular Solids
3.2.6 Metallic Solids and Alloys
3.2.7 Noncovalent Interactions in Biomolecules and Polymers
3.3.1 Characteristics of Solids
3.3.2 Characteristics of Liquids
3.3.3 Transition Between Solid and Liquid Phases
3.3.4 Characteristics of Gases
3.3.5 Comparing the Three States of Matter
3.4.1 Understanding the Ideal Gas Law
3.4.2 Partial Pressures in Gas Mixtures
3.4.3 Graphical Representations of Gas Laws
3.4.4 Applications of the Ideal Gas Law
3.4.5 Limitations of the Ideal Gas Law
3.5.1 Fundamentals of Kinetic Molecular Theory
3.5.2 Maxwell-Boltzmann Distribution
3.5.3 Relationship Between Kinetic Energy, Velocity, and Temperature
3.5.4 Graphical Representations of Particle Motion
3.5.5 Applications of KMT to Gas Properties
3.6.1 Reasons for Deviation from Ideal Gas Law
3.6.2 Interparticle Attractions and Deviations
3.6.3 Impact of Particle Volume on Gas Behavior
3.6.4 Correcting for Non-Ideal Behavior: Van der Waals Equation
3.6.5 Graphical Representation of Deviations
3.7.1 Types of Mixtures: Solutions vs. Heterogeneous Mixtures
3.7.2 Composition of Solutions
3.7.3 Expressing Solution Concentration: Molarity
3.7.4 Calculating the Number of Solute Particles and Volume
3.7.5 Applications of Molarity in the Laboratory
3.8.1 Basics of Particulate Models
3.8.2 Representing Interactions Between Components
3.8.3 Depicting Concentrations in Particulate Models
3.8.4 Applications of Particulate Models
3.8.5 Limitations of Particulate Models
3.9.1 Principles of Chromatography
3.9.2 Types of Chromatography
3.9.3 Intermolecular Interactions in Chromatography
3.9.4 Practical Applications of Chromatography
3.9.5 Analyzing Chromatographic Data
3.10.1 Principles of Solubility
3.10.2 Solubility of Ionic and Molecular Compounds
3.10.3 Factors Affecting Solubility
3.10.4 Solubility and Chemical Reactions
3.10.5 Measuring and Expressing Solubility
3.11.1 Overview of the Electromagnetic Spectrum
3.11.2 Microwave Radiation and Rotational Transitions
3.11.3 Infrared Radiation and Vibrational Transitions
3.11.4 Ultraviolet/Visible Radiation and Electronic Transitions
3.11.5 Applications of Spectroscopy in Chemistry
3.11.6 Analyzing Spectroscopic Data
3.12.1 Basics of the Photoelectric Effect
3.12.2 Introduction to Photon Absorption and Electronic Transitions
3.12.3 Relationship Between Wavelength, Frequency, and Energy
3.12.4 Applications of the Photoelectric Effect
3.12.5 Analyzing Energy Changes in Electronic Transitions
3.13.1 Fundamentals of the Beer-Lambert Law
3.13.2 Molar Absorptivity (E)
3.13.3 Impact of Path Length and Concentration
3.13.4 Practical Applications of the Beer-Lambert Law
3.13.5 Limitations and Considerations of the Beer-Lambert Law
4.1.1 Distinguishing Between Physical and Chemical Changes
4.1.2 Introduction to Physical Changes
4.1.3 Evidence of Chemical Changes
4.1.4 Representing Chemical Changes with Equations
4.1.5 Importance of Identifying Chemical Reactions
4.2.1 Principles of Chemical Equations
4.2.2 Types of Chemical Equations
4.2.3 Writing Complete Ionic Equations
4.2.4 Writing Net Ionic Equations
4.2.5 Identifying Spectator Ions
4.2.6 Applications and Importance of Net Ionic Equations
4.3.1 Introduction to Particulate Models
4.3.2 Translating Chemical Equations to Particulate Models
4.3.3 Types of Chemical Reactions and Their Particulate Representations
4.3.4 Physical Changes in Particulate Models
4.3.5 Challenges in Representing Complex Reactions
4.3.6 Educational Value of Particulate Models
4.4.1 Classification of Processes
4.4.2 Chemical Processes and Bond Interactions
4.4.3 Physical Processes and Intermolecular Interactions
4.4.4 Grey Areas in Classification: Dissolution of Salts
4.4.5 Implications of Process Classification
4.4.6 Educational Approaches to Teaching Changes
4.5.1 Fundamentals of Stoichiometry
4.5.2 Reading and Interpreting Chemical Equations
4.5.3 Stoichiometric Calculations: Reactants and Products
4.5.4 Stoichiometry and the Mole Concept
4.5.5 Application of Stoichiometry in Gases and Solutions
4.5.6 Limiting Reactants and Excess Reactants
4.5.7 Percent Yield and Theoretical Yield
4.6.1 Basics of Titration
4.6.2 The Equivalence Point
4.6.3 Indicators and the Endpoint
4.6.4 Calculating Concentrations from Titration Data
4.6.5 Types of Titration
4.6.6 Practical Considerations in Titration
4.7.1 Acid-Base Reactions
4.7.2 Oxidation-Reduction Reactions
4.7.3 Identifying Oxidized and Reduced Species
4.7.4 Precipitation Reactions
4.7.5 Applications of Reaction Types
4.7.6 Experimentally Determining Reaction Types
4.8.1 Brønsted-Lowry Acid-Base Theory
4.8.2 Role of Water in Acid-Base Reactions
4.8.3 Identifying Conjugate Acid-Base Pairs
4.8.4 Relative Strengths of Acids and Bases
4.8.5 Applications of Acid-Base Reactions
4.8.6 Experimentally Identifying Acids and Bases
4.9.1 Fundamentals of Redox Reactions
4.9.2 Constructing Redox Equations from Half-Reactions
4.9.3 Identifying Oxidation and Reduction Half-Reactions
4.9.4 Balancing Redox Equations
4.9.5 Applications of Redox Reactions
4.9.6 Experimental Determination of Redox Reactions
5.1.1 Introduction to Reaction Kinetics
5.1.2 Measuring Reaction Rates
5.1.3 Factors Influencing Reaction Rates
5.1.4 The Effect of Concentration on Reaction Rates
5.1.5 The Role of Temperature in Reaction Rates
5.1.6 Catalysts and Reaction Rates
5.1.7 Surface Area and Reaction Rates
5.2.1 Understanding Rate Law
5.2.2 Reaction Order
5.2.3 Rate Constant
5.2.4 Determining Rate Laws from Experimental Data
5.3.1 Graphical Methods for Determining Reaction Order
5.3.2 Determining First Order Reactions
5.3.3 Determining Second Order Reactions
5.3.4 Calculating the Rate Constant
5.3.5 Understanding Half-Life in First Order Reactions
5.3.6 First Order Kinetics and Radioactive Decay
5.4.1 Understanding Elementary Reactions
5.4.2 Rate Laws for Elementary Reactions
5.4.3 Molecularity of Elementary Reactions
5.4.4 Representing Rate Laws from Stoichiometry
5.4.5 Limitations in Elementary Reaction Collisions
5.5.1 Basics of the Collision Model
5.5.2 Criteria for Successful Collisions
5.5.3 Impact of Collision Frequency on Reaction Rate
5.5.4 Maxwell-Boltzmann Energy Distribution
5.5.5 Temperature Dependence of Reaction Rates
5.6.1 Basics of Reaction Energy Profiles
5.6.2 Bonds and Elementary Reactions
5.6.3 The Reaction Coordinate and Transition State
5.6.4 Activation Energy
5.6.5 The Arrhenius Equation and Activation Energy
5.6.6 Interpreting Reaction Energy Profiles
5.7.1 Basics of Reaction Mechanisms
5.7.2 Elementary Steps and the Overall Reaction
5.7.3 Reaction Intermediates
5.7.4 Experimental Evidence for Mechanisms
5.7.5 Distinguishing Reaction Mechanisms
5.7.6 Significance of Reaction Mechanisms
5.8.1 Understanding Rate-Limiting Steps
5.8.2 Rate Law and Reaction Mechanisms
5.8.3 Identifying Rate-Limiting Steps
5.8.4 Examples of Rate Laws from Mechanisms
5.8.5 Irreversible and Reversible Elementary Steps
5.9.1 Fundamentals of Steady-State Approximation
5.9.2 Applying Steady-State Approximation to Reaction Mechanisms
5.9.3 Deriving Rate Laws with Steady-State Approximation
5.9.4 Limitations of the Steady-State Approximation
5.10.1 Basics of Multistep Reaction Energy Profiles
5.10.2 Constructing Energy Profiles for Multistep Reactions
5.10.3 Activation Energy in Multistep Reactions
5.10.4 Interpreting Energy Profiles
5.10.5 Energetics and Reaction Mechanisms
5.11.1 Role of Catalysts in Chemical Reactions
5.11.2 Mechanism of Catalysis
5.11.3 Types of Catalysis
5.11.4 Catalysis and Reaction Intermediates
5.11.5 Examples of Catalytic Mechanisms
5.11.6 Impact of Catalysis on Reaction Rates
6.1.1 Energy Changes and Temperature
6.1.2 Defining Endothermic and Exothermic Processes
6.1.3 Energy Flow in Chemical Reactions
6.1.4 Solution Formation and Energy Changes
6.1.5 Experimental Observation of Energy Changes
6.2.1 Introduction to Energy Diagrams
6.2.2 Constructing Energy Diagrams
6.2.3 Interpreting Exothermic and Endothermic Processes
6.2.4 Activation Energy and Reaction Pathways
6.2.5 Applications of Energy Diagrams
6.3.1 Thermal Energy and Molecular Motion
6.3.2 Mechanisms of Heat Transfer
6.3.3 Achieving Thermal Equilibrium
6.3.4 Principles of Heat Exchange
6.3.5 Practical Implications of Thermal Equilibrium
6.4.1 Fundamentals of Heat Transfer
6.4.2 First Law of Thermodynamics
6.4.3 Specific Heat Capacity
6.4.4 Calorimetry Experiments
6.4.5 Calculating Heat Changes
6.4.6 Energy Changes in Chemical Systems
6.5.1 Thermodynamics of Phase Changes
6.5.2 Energy and Phase Transitions
6.5.3 Molar Enthalpy of Phase Changes
6.5.4 Calculating Energy Changes in Phase Transitions
6.5.5 Conservation of Energy in Phase Changes
6.6.1 Enthalpy Change of a Reaction
6.6.2 Calculating Enthalpy Change
6.7.1 Role of Bond Breaking and Formation
6.7.2 Calculating Reaction Enthalpy from Bond Energies
6.7.3 Understanding Average Bond Enthalpies
6.7.4 Exothermic vs. Endothermic Reactions
6.8.1 Standard Enthalpies of Formation
6.8.2 Calculating Standard Enthalpies of Reactions
6.9.1 Principles of Hess's Law
6.9.2 Enthalpy and Reaction Sequences
6.9.3 Thermal Energy Transfer and Equilibrium
7.1 Introduction to Equilibrium
7.2.1 Relationship Between Reaction Direction and Rates
7.2.2 Establishing Equilibrium
7.3.1 Understanding the Reaction Quotient (Q)
7.3.3 Equilibrium Expressions for Kc and Kp
7.4.1 Determining Equilibrium Constants from Experiments
7.5.1 Significance of K Values
7.5.2 Relationship Between K and Concentrations at Equilibrium
7.6.1 Inversion of K for Reversed Reactions
7.6.2 Effect of Stoichiometric Changes
7.6.3 Summation of Reactions and K
7.6.4 Algebraic Manipulations of K and Q
7.7.1 Predicting Equilibrium Concentrations
7.8.1 Particulate Models of Equilibrium
7.8.2 Connecting Particulate Models to Equilibrium Constants
7.9.1 Understanding Le Châtelier’s Principle
7.9.2 Application of Le Châtelier’s Principle
7.10.1 Relationship Between Q, K, and Equilibrium
7.10.2 Effects of Stresses on Q and K
7.11.1 Solubility Equilibria and Ksp
7.11.2 Calculating Solubility from Ksp
7.11.3 Predicting Solubility of Substances
7.11.4 Relationship Between Ksp and Solubility Rules
7.12.1 Understanding the Common-Ion Effect
7.12.2 Common-Ion Effect and Le Châtelier’s Principle
7.12.3 Calculating Solubility with Common Ions
7.13.1 pH Sensitivity of Salt Solubility
7.13.2 Applying Le Châtelier’s Principle to pH and Solubility
7.14.1 Relationship Between Solubility and Free Energy Changes
7.14.2 Enthalpic and Entropic Contributions
8.1.1 Understanding pH and pOH
8.1.2 Autoionization of Water
8.1.3 Neutral Solutions and Water
8.1.4 Temperature Dependence of Kw
8.2.1 pH of Strong Acid Solutions
8.2.2 pOH and pH of Strong Base Solutions
8.3.1 Partial Ionization of Weak Acids
8.3.2 Equilibrium in Weak Acid Solutions
8.3.3 Partial Ionization of Weak Bases
8.3.4 Equilibrium in Weak Base Solutions
8.3.5 Percent Ionization of Weak Acids and Bases
8.4.1 Reactions of Strong Acids and Bases
8.4.2 Reactions of Weak Acids with Strong Bases
8.4.3 Reactions of Weak Bases with Strong Acids
8.4.4 Reactions of Weak Acids with Weak Bases
8.5.1 Overview of Acid-Base Titrations
8.5.2 Equivalence Point in Titrations
8.5.3 Half-Equivalence Point and pKa Determination
8.5.4 Titrations of Polyprotic Acids
8.6.1 Acid Strength and Molecular Structure
8.6.2 Characteristics of Strong Acids
8.6.3 Weak Acids and Bases
8.6.4 Strong Bases and Their Conjugate Acids
8.6.5 Effect of Electronegativity on Acid Strength
8.7.1 Relationship Between pH and pKa
8.7.2 Use of Acid-Base Indicators
8.8.1 Buffer Solution Composition
8.8.2 Mechanism of pH Stabilization by Buffers
8.9.1 Relationship Between pH and Buffer Components
8.9.2 Impact of Adding Acid or Base to a Buffer
8.10.1 Role of Concentration in Buffer Capacity
8.10.2 Buffer Composition and Capacity
9.1.1 Entropy and Matter Dispersion
9.1.2 Entropy in Gas-Phase Reactions
9.1.3 Entropy and Energy Dispersion
9.2.1 Calculating Entropy Change
9.3.1 Understanding Gibbs Free Energy (ΔGo)
9.3.2 Thermodynamic Favorability and ΔGo
9.3.3 Calculating ΔGo from ΔGf°
9.3.4 Role of Enthalpy and Entropy in Determining ΔGo
9.3.5 Predicting Thermodynamic Favorability from ΔHo and ΔSo
9.4.1 Thermodynamic Favorability vs. Kinetic Feasibility
9.4.2 Kinetic Control and Activation Energy
9.4.3 Implications of Kinetic Control
9.5.1 Thermodynamic Favorability and Equilibrium
9.5.2 Relationship Between K and ΔG°
9.5.3 Qualitative Connections Between K and ΔG°
9.5.4 ΔG° and the Direction of Equilibrium
9.6.1 Use of External Energy Sources
9.6.2 Mechanism of Coupled Reactions
9.6.3 Examples of Coupled Reactions in Biological Systems
9.7.1 Components of Electrochemical Cells
9.7.2 Operational Principles of Galvanic vs. Electrolytic Cells
9.7.3 Oxidation and Reduction in Electrochemical Cells
9.8.1 Thermodynamics of Electrochemical Cells
9.8.2 Calculating Standard Cell Potential
9.8.3 Relationship Between ΔG° and E°
9.9.1 Cell Potential and Reaction Equilibrium
9.9.2 Impact of Nonstandard Conditions on Cell Potential
9.9.3 Cell Potential in Concentration Cells
9.9.4 Qualitative Use of the Nernst Equation
9.10.1 Principles of Electrolysis
9.10.2 Faraday's Law and Redox Stoichiometry
9.10.3 Calculating Mass Changes at Electrodes
9.10.4 relating Current and Charge Flow

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