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AP Chemistry

Study Notes
1.1.1 Why Chemists Use the Mole
1.1.2 Avogadro’s Number and Counting Particles
1.1.3 Molar Mass and Atomic Mass Units
1.1.4 Dimensional Analysis with n = m/M
1.2.1 What a Mass Spectrum Tells You
1.2.2 Calculating Average Atomic Mass
1.2.3 What You Don’t Need to Interpret
1.3.1 Pure Substances: Molecules vs Formula Units
1.3.2 Law of Definite Proportions
1.3.3 Empirical Formula Basics
1.3.4 From Percent Composition to Empirical Formula
1.4.1 Pure Substances vs Mixtures
1.4.2 Elemental Analysis and Purity
1.4.3 Finding the Composition of a Mixture
1.5.1 Subatomic Particles and Charge
1.5.2 Coulomb’s Law in Atomic Interactions
1.5.3 Shells, Subshells, and Orbital Filling
1.5.4 Aufbau Principle: Writing Electron Configurations
1.5.5 Electron Configurations of Ions
1.5.6 Core vs Valence Electrons
1.5.7 Ionization Energy, Shielding, and Effective Nuclear Charge
1.5.8 What Is Excluded on the AP Exam
1.6.1 PES: Measuring Electron Energies
1.6.2 Reading Peaks: Binding Energy and Electron Count
1.6.3 Connecting PES to Electron Configuration and Nuclear Attraction
1.7.1 Why the Periodic Table Is Organized the Way It Is
1.7.2 Shielding and Effective Nuclear Charge (Zeff)
1.7.3 Ionization Energy Trend
1.7.4 Atomic and Ionic Radius Trends
1.7.5 Electron Affinity and Electronegativity Trends
1.7.6 Using Periodicity to Estimate Missing Data
1.7.7 What Is Excluded on the AP Exam
1.8.1 Valence Electrons and Bond Formation
1.8.2 Families Form Similar Compounds
1.8.3 Predicting Ionic Charges from the Periodic Table
1.8.4 Reactivity Trends and Periodicity
2.1.1 Electronegativity trends and bonding predictions
2.1.2 Nonpolar covalent bonds from similar electronegativity
2.1.3 Polar covalent bonds, partial charges, and dipoles
2.1.4 Ionic–covalent continuum and classifying bonds
2.1.5 Metallic bonding and delocalized electrons
2.2.1 Potential energy curves: bond length and bond energy
2.2.2 Covalent bond length and bond order effects
2.2.3 Coulomb’s law for ionic interaction strength
2.3.1 Ionic crystal lattices and macroscopic properties
2.4.1 Metallic solids: ions in a sea of electrons
2.4.2 Interstitial alloys and their structure–property links
2.4.3 Substitutional alloys and lattice substitution
2.5.1 Constructing Lewis diagrams using core principles
2.6.1 Resonance as a refinement to Lewis structures
2.6.2 Using the octet rule and formal charge to choose a model
2.6.3 Limits of Lewis structures, including odd-electron species
2.7.1 VSEPR theory: electron-pair repulsion and shapes
2.7.2 From Lewis + VSEPR to key molecular properties
2.7.3 Hybridization nomenclature and ideal bond angles
2.7.4 Sigma and pi bonds from orbital overlap
2.7.5 Restricted rotation and geometric isomerism in pi bonds
2.7.6 What you are responsible for on the AP Exam (exclusions)
3.1.1 London Dispersion Forces (LDF) and Temporary Dipoles
3.1.2 LDF vs Van der Waals Terminology
3.1.3 Dipole Moments and Dipole–Dipole Interactions
3.1.4 Ion–Dipole Forces and Relative Strengths
3.1.5 Comparing Dipole–Dipole and Ion–Dipole Using Partial Charges
3.1.6 Hydrogen Bonding: When and Why It Is Strong
3.1.7 Noncovalent Interactions in Large Biomolecules
3.2.1 Intermolecular Forces and Phase-Change Properties
3.2.2 Using Particulate Representations to Explain Properties
3.2.3 Ionic Solids: Melting Point, Brittleness, and Conductivity
3.2.4 Covalent Network Solids: Diamond, Graphite, and High Melting Points
3.2.5 Molecular Solids and Weak Intermolecular Attractions
3.2.6 Metallic Solids and Alloys: The Sea of Electrons
3.2.7 Biomolecules and Polymers: Shape from Noncovalent Forces
3.3.1 Particulate Models of Phases
3.3.2 Crystalline vs Amorphous Solids
3.3.3 Packing and Interparticle Interactions in Solids
3.3.4 Liquids: Close Contact, Motion, and Forces
3.3.5 Why Solids and Liquids Have Similar Molar Volumes
3.3.6 Gases: Motion, Spacing, and State Variables
3.3.7 Phase Diagrams Not Assessed
3.4.1 Using PV = nRT to Relate Gas Properties
3.4.2 Dalton’s Law and Mole Fraction
3.4.3 Graphs of P–V–T–n Relationships
3.5.1 KMT Links Particle Motion to Gas Properties
3.5.2 Kinetic Energy and Molecular Speed
3.5.3 Temperature in Kelvin and Average Kinetic Energy
3.5.4 Interpreting Maxwell–Boltzmann Distributions
3.6.1 Why Real Gases Deviate from PV = nRT
3.6.2 Attractions and Condensation Conditions
3.6.3 Finite Particle Volume at High Pressure
3.7.1 Solutions vs Heterogeneous Mixtures
3.7.2 States of Solutions and Particle Counting
3.7.3 Molarity Calculations (M = n/L)
3.8.1 Showing Interactions in Particulate Models
3.8.2 Showing Concentration in Particulate Models
3.8.3 Excluded: Colligative Properties and Alternative Concentration Units
3.9.1 Why Filtration Fails for Liquid Solutions
3.9.2 Separation Based on Intermolecular Interactions
3.9.3 Chromatography: Mobile vs Stationary Phase Interactions
3.9.4 Using a Chromatogram to Infer Polarity
3.10.1 Like Dissolves Like: Intermolecular Similarity
3.10.2 Solubility of Ionic and Molecular Compounds in Different Solvents
3.11.1 Connecting Spectrum Regions to Transitions
3.11.2 Microwave Radiation and Molecular Rotation
3.11.3 Infrared Radiation and Molecular Vibration
3.11.4 Ultraviolet/Visible Radiation and Electronic Transitions
3.12.1 Energy Change on Absorption or Emission
3.12.2 Wave Relationship: c = λν
3.12.3 Planck Relationship: E = hν
3.13.1 Beer–Lambert Equation and Variable Meanings
3.13.2 Why Absorbance Is Proportional to Concentration
3.13.3 Choosing λmax and Using a Spectrophotometer
3.13.4 Experimental Error Considerations in Spectroscopy
4.1.1 Physical changes vs chemical composition
4.1.2 Evidence of chemical change
4.2.1 Representing changes with balanced equations
4.2.2 Conservation in chemical equations
4.2.3 Molecular, complete ionic, and net ionic forms
4.2.4 Writing and balancing net ionic equations
4.3.1 Particulate models from balanced equations
4.3.2 Connecting symbols to particles and quantities
4.4.1 Bond changes vs intermolecular changes
4.4.2 Borderline cases: dissolving ionic solids
4.5.1 Using conservation to relate reactants and products
4.5.2 Mole ratios from coefficients
4.5.3 Stoichiometry with gases and solutions
4.5.4 Limiting and excess reactants (application)
4.6.1 Purpose of titration and key terms
4.6.2 Equivalence point vs endpoint
4.6.3 Using stoichiometry at equivalence
4.7.1 Identifying acid–base reactions
4.7.2 Identifying redox reactions and combustion
4.7.3 Oxidation vs reduction and electron flow
4.7.4 Using oxidation numbers to find species changed
4.7.5 Identifying precipitation reactions and key solubilities
4.7.6 What you don’t need to memorize (exam note)
4.8.1 Brønsted–Lowry acids, bases, and conjugate pairs
4.8.2 Water’s role in aqueous acid–base reactions
4.8.3 Comparing strengths using conjugate pairs
4.8.4 Scope for AP Chemistry (exam note)
4.9.1 Balancing redox with half-reactions
5.1.1 What Is Reaction Rate (Kinetics Basics)
5.1.2 Relating Reactant and Product Rates Using Stoichiometry
5.1.3 Factors That Affect Reaction Rate
5.2.1 Measuring Reaction Rate Experimentally
5.2.2 Writing Rate Laws (Rate vs. Concentration)
5.2.3 Reaction Order and Overall Order
5.2.4 Rate Constant k: Meaning, Units, and Temperature Dependence
5.2.5 The Method of Initial Rates
5.3.1 Determining Reaction Order from Concentration–Time Data
5.3.2 Zero-Order Integrated Rate Law and Linear Plots
5.3.3 First-Order Integrated Rate Law and Half-Life
5.3.4 Second-Order Integrated Rate Law and Reciprocal Plots
5.3.5 Finding k from the Slope of Linearized Plots
5.3.6 Radioactive Decay as a First-Order Process
5.4.1 Molecularity and Rate Laws for Elementary Steps
5.4.2 Why Termolecular Elementary Steps Are Rare
5.5.1 Collision Theory and Effective Collisions
5.5.2 Activation Energy and Orientation Requirements
5.5.3 Maxwell–Boltzmann Distribution and Temperature Effects
5.6.1 Reaction Coordinate, Bonds, and the Transition State
5.6.2 Reading an Energy Profile: Ea and Overall Energy Change
5.6.3 Why Temperature Affects Rate (Transition State Idea)
5.6.4 Arrhenius Equation (Conceptual, No Calculations)
5.7.1 What Is a Reaction Mechanism?
5.7.2 Checking a Mechanism Against the Overall Equation
5.7.3 Identifying Reaction Intermediates (and Catalysts)
5.7.4 Evidence for Mechanisms: Detecting Intermediates
5.8.1 Rate-Determining Step and the Rate Law
5.8.2 Using Molecularity of the Slow Step
5.9.1 When the First Step Is Not Rate Limiting
5.9.2 Pre-Equilibrium: Getting a Rate Law from a Mechanism
5.10.1 Constructing Multistep Energy Profiles from Step Energetics
5.10.2 Intermediates and Transition States on Multistep Profiles
5.10.3 Linking the Largest Barrier to the Slow Step (Conceptual)
5.11.1 How Catalysts Speed Up Reactions
5.11.2 Catalysts Are Regenerated (Net Amount Constant)
5.11.3 Catalysis by Binding and Orientation (Enzymes)
5.11.4 Covalent and Acid–Base Catalysis
5.11.5 Surface Catalysis
6.1.1 Temperature Change as Evidence of Energy Change
6.1.2 Classifying Processes: Endothermic vs. Exothermic
6.1.3 Energy Flow in Chemical Reactions (System vs. Surroundings)
6.1.4 Solution Formation: Interparticle Forces and Energy
6.2.1 Reading Energy Diagrams for Endothermic and Exothermic Changes
6.2.2 Sketching Energy Diagrams with Correct Labels and Units
6.3.1 Kinetic Molecular View: Why Warmer Means Faster Particles
6.3.2 How Collisions Cause Heat Transfer
6.3.3 Thermal Equilibrium: When Temperatures Become Equal
6.4.1 Quantifying Heat Transfer with q = mcΔT
6.4.2 Calorimetry: Measuring Heat in the Lab
6.4.3 First Law of Thermodynamics in Calorimetry
6.4.4 Why Different Materials Warm Up Differently
6.4.5 Heating vs. Cooling and System Energy
6.4.6 Specific Heat Capacity vs. Molar Heat Capacity
6.4.7 Energy Changes in Chemical Systems: Three Pathways
6.4.8 Dissolution Calorimetry: Using Temperature to Infer Energy Flow
6.5.1 Melting and Boiling: Energy Increases at Constant Temperature
6.5.2 Freezing and Condensing: Energy Decreases at Constant Temperature
6.5.3 Complementary Phase Changes and Sign Conventions
6.5.4 Using Molar Enthalpies (Fusion and Vaporization) in Calculations
6.6.1 Interpreting ΔH at Constant Pressure
6.6.2 Exothermic vs. Endothermic Reactions and Surroundings
6.6.3 Why Reactions Change Temperature: Bonds and Energy
6.6.4 Calculating Reaction Heat from Moles and Molar Enthalpy
6.7.1 Bond Breaking and Bond Forming Change Potential Energy
6.7.2 Estimating Energy In: Sum of Bonds Broken
6.7.3 Estimating Energy Out: Sum of Bonds Formed
6.7.4 Deciding Endothermic vs. Exothermic from Bond Energies
6.8.1 What Standard Enthalpy of Formation Tables Tell You
6.8.2 Using ΔH°reaction = ΣΔHf°products − ΣΔHf°reactants
6.9.1 Breaking a Process into Steps with Individual Energy Changes
6.9.2 Why Enthalpy Changes Add: Energy Conservation
6.9.3 Reversing Reactions: Same Magnitude, Opposite Sign
6.9.4 Scaling Reactions: Enthalpy Scales by the Same Factor
6.9.5 Adding Reactions: Add Enthalpies to Get the Overall ΔH
7.1.1 Reversible processes: physical and chemical
7.1.2 Macroscopic signs of equilibrium
7.1.3 Dynamic equilibrium and equal rates
7.1.4 Reading equilibrium graphs
7.2.1 Predicting net direction from rates
7.2.2 When equilibrium is established
7.3.1 Writing Q and K expressions (Qc/Qp, Kc/Kp)
7.3.2 How Q compares to K
7.3.3 What to omit from Q/K (and exam notes)
7.4.1 Determining K from equilibrium data
7.5.1 Large K: product-favored equilibria
7.5.2 Small K: reactant-favored equilibria
7.6.1 Reversing a reaction inverts K
7.6.2 Scaling a reaction changes K by a power
7.6.3 Adding reactions multiplies K values
7.6.4 K and Q share the same algebra
7.7.1 Predicting equilibrium amounts from initial conditions
7.7.2 Using Q vs K to decide the shift direction
7.7.3 Recognizing equilibrium (Q = K)
7.8.1 Particulate models before and at equilibrium
7.8.2 Linking particle ratios to K
7.9.1 Predicting responses to common stresses
7.9.2 Connecting shifts to measurable changes
7.10.1 Disturbances make Q differ from K
7.10.2 Re-establishing a new equilibrium
7.10.3 Which stresses change Q vs change K
7.11.1 Dissolution as an equilibrium process (Ksp)
7.11.2 Calculating and comparing solubility using Ksp
7.11.3 Relating solubility rules to Ksp
7.11.4 Using molar solubility to find Ksp
7.12.1 Explaining the common-ion effect qualitatively
7.12.2 Calculating common-ion effects with Ksp
8.1.1 pH and pOH as measures of [H3O+] and [OH−]
8.1.2 Hydronium vs hydrogen ion notation
8.1.3 Water autoionization and Kw
8.1.4 Neutral water, pKw, and temperature effects
8.2.1 Strong acids: complete ionization and pH
8.2.2 Strong bases: dissociation and pOH
8.2.3 Converting between pH and pOH using pKw
8.3.1 Weak acids: partial ionization and equilibrium picture
8.3.2 Ka and pKa: finding pH of a weak acid
8.3.3 Weak bases and Kb / pKb: finding pH of a weak base
8.3.4 Percent ionization of weak acids and bases
8.3.5 Linking conjugate pairs with Kw
8.4.1 Strong acid–strong base mixtures
8.4.2 Weak acid + strong base: excess acid, excess base, or equimolar
8.4.3 Weak base + strong acid: excess base, excess acid, or equimolar
8.4.4 Weak acid + weak base mixtures
8.5.1 Titrations and titration curves
8.5.2 Equivalence point and finding analyte concentration
8.5.3 Half-equivalence point and determining pKa
8.5.4 What sets pH at equivalence for strong vs weak systems
8.5.5 Polyprotic titrations: number of protons and species (qualitative)
8.6.1 Identifying acidic protons from structure
8.6.2 Why some acids are stronger: stabilizing conjugate bases
8.6.3 Common structural examples of weak/strong acids and bases
8.6.4 Electronegativity and acid strength trends
8.7.1 Predicting predominant form from pH vs pKa/pKb
8.7.2 How acid–base indicators work
8.7.3 Choosing an indicator for a titration endpoint
8.8.1 What makes a solution a buffer
8.8.2 How buffers stabilize pH when acid or base is added
8.9.1 Using Henderson–Hasselbalch to find buffer pH
8.9.2 Why small additions cause small pH changes in a buffer
8.9.3 Exam scope for Henderson–Hasselbalch applications
8.10.1 Concentration vs capacity at constant ratio
8.10.2 Which component is in excess determines capacity direction
8.11.1 When solubility depends on pH
8.11.2 Predicting pH effects on solubility using Le Châtelier
8.11.3 Exam scope for pH–solubility questions
9.1.1 What Entropy Measures: Dispersal and Disorder
9.1.2 Predicting ΔS for Phase Changes, Volume Changes, and Gas Moles
9.1.3 Temperature and Energy Dispersal (KMT and Kinetic Energy Distributions)
9.2.1 Absolute (Standard Molar) Entropy and Standard Conditions
9.2.2 Calculating ΔS°rxn Using Stoichiometry
9.2.3 Interpreting Results: Units, Sign, and Common Pitfalls
9.3.1 Standard States and the Meaning of ΔG°
9.3.2 Thermodynamically Favored vs. Unfavored (ΔG° Sign)
9.3.3 Finding ΔG°rxn from Standard Free Energies of Formation
9.3.4 When Enthalpy and Entropy Must Both Be Considered
9.3.5 Using ΔG° = ΔH° − TΔS° at a Given Temperature
9.3.6 Predicting Temperature Conditions for Favorability from Signs of ΔH° and ΔS°
9.4.1 Thermodynamically Favored Does Not Mean Fast
9.4.2 Kinetic Control and Activation Energy Barriers
9.4.3 Kinetics vs. Equilibrium: What Slow Rates Do (and Don’t) Imply
9.4.4 Reasoning from Evidence: Identifying Kinetic Control in Practice
9.5.1 Connecting ΔG° and K Under Standard Conditions
9.5.2 The Mathematical Relationship: ΔG° and K
9.5.3 Estimating K from the Size of ΔG° Relative to RT
9.5.4 Interpreting Signs: Product-Favored vs Reactant-Favored
9.6.1 What ΔG° of Dissolution Represents
9.6.2 Enthalpy and Entropy Contributions: Breaking the Solid Apart
9.6.3 Solvent Reorganization (Hydration/Solvation Shells)
9.6.4 Solute–Solvent Interactions and Their Thermodynamic Impact
9.6.5 Why Solubility Predictions Can Be Hard: Competing Terms and Cancellations
9.7.1 Driving Unfavorable Processes with External Electrical Energy
9.7.2 Driving Unfavorable Processes with Light Energy
9.7.3 Reaction Coupling: Using a Favorable Reaction to Power an Unfavorable One
9.7.4 Net Favorability: Adding Reactions to Obtain Overall ΔG° < 0
9.8.1 Core Components of Electrochemical Cells and Their Roles
9.8.2 Tracking Electron Flow, Ion Flow, and Observable Changes
9.8.3 Galvanic vs. Electrolytic Cells: Favored vs. Unfavored Reactions
9.8.4 Using Cell Diagrams to Locate Half-Reactions and Current Direction
9.8.5 Anode vs. Cathode: Where Oxidation and Reduction Occur
9.9.1 Redox Reactions in Cells and What Voltage Sign Means
9.9.2 Calculating E°cell from Standard Reduction Potentials
9.9.3 Linking Cell Potential to Free Energy (ΔG° = −nFE°)
9.10.1 How Concentration Changes Affect Cell Potential
9.10.2 Q, K, and the Point Where E Reaches Zero
9.10.3 Why Le Châtelier’s Principle Doesn’t Apply to Cells
9.10.4 Concentration Cells and Direction of Spontaneous Electron Flow
9.10.5 Qualitative Use of the Nernst Equation (Concept Over Plug-and-Chug)
9.11.1 Faraday’s Law Basics: Current, Charge, and Time
9.11.2 Electrons and Stoichiometry in Electrolysis Problems
9.11.3 Mass Changes at Electrodes: Electroplating and Dissolution
9.11.4 Solving for Current or Time from Amount of Substance Produced
9.11.5 Accounting for Ion Charge and Multiple-Electron Transfers

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