3rd CHEM REVIEWER PDF

CHAPTER 1: INTERMOLECULAR FORCES & PROPERTIES OF LIQUIDS AND SOLIDS KINETIC MOLECULAR THEORY (PART 1) Fundamental Principles Core Concepts: ​ All matter consists of tiny particles ​ Particles are in constant motion ​ Particle speed is directly proportional to temperature ​ Different sta...

CHAPTER 1: INTERMOLECULAR FORCES & PROPERTIES OF LIQUIDS AND SOLIDS KINETIC MOLECULAR THEORY (PART 1) Fundamental Principles Core Concepts: ​ All matter consists of tiny particles ​ Particles are in constant motion ​ Particle speed is directly proportional to temperature ​ Different states of matter vary in: 1.​ Particle distance 2.​ Particle motion freedom 3.​ Particle interaction extent Phases of Matter Comparison Property Solid Liquid Gas Particle Distance Closest Intermediate Furthest Particle Movement Minimal Moderate Maximum Empty Space Least Intermediate Most Temperature Effects Molecular Behavior: ​ Increased temperature - Increased molecular motion ​ Molecules gain kinetic energy ​ Potential phase transitions occur INTERMOLECULAR FORCES (PART 2) Definition: Intermolecular Forces ​ Attractive forces between molecules ​ Significantly weaker than intermolecular bonds ​ Primarily active in solid and liquid states Types of Intermolecular Forces 1. London Dispersion Forces (LDF) Characteristics: ​ Present in all molecules ​ Result from temporary electron distribution  ​ Strength depends on: - Number of electrons - Electron cloud size - Molecular complexity 2. Dipole-Dipole Forces (DDF) Key Features: ​ Occur between polar molecules ​ Electrostatic attraction ​ Strength proportional to dipole moment 𝑞1𝑞2 ​ Follows Coulomb’s law: F= 2 4πϵ0𝑟 3. Ion-Dipole Forces (IDF) Interaction: ​ Between ions and polar molecules ​ Strength determined by: - Ion charge magnitude - Molecule’s dipole moment 4. Hydrogen Bonding (H-BOND) Unique Characteristics: ​ Strongest intermolecular force ​ Occurs between H and O, N, F atoms ​ Particularly strong in water molecules Force Strength Ranking 1.​ Hydrogen Bonds (Strongest) 2.​ Ion-Dipole Forces 3.​ Dipole-Dipole Forces 4.​ London Dispersion Forces (Weakest) PROPERTIES OF LIQUIDS & INTERMOLECULAR FORCES (PART 3) Fundamental Concepts Fluid Definition: ​ Fluid: A substance that can flow (includes both gasses and liquids) ​ Characterized by ability to move and change shape. Key Properties of Liquids Surface Tension (Intermolecular Forces Impact) Definition: Measure of elastic force at liquid’s surface Key Characteristics: ​ Energy required to stretch liquid surface ​ Molecules at surface pulled downward and sideways ​ Stronger intermolecular forces = Higher surface tension Viscosity (Resistance to Flow) Definition: Fluid’s resistance to flowing Measurement: ​ Time taken to flow through vertical tube ​ Expressed in centipoise ​ Water = 1 centipoise at 20 degree Factors Affecting Viscosity: ​ Intermolecular Forces ​ Molecular Structure - More -OH groups = Stronger hydrogen bonding - Large molecules = Greater viscosity Vapor Pressure Definition: Pressure exerted by gas in equilibrium with liquid Key Insights: ​ Depends on intermolecular force strength ​ Weak intermolecular forces = Higher vapor pressure ​ Strong intermolecular forces = Lower vapor pressure Molar Heat of Vaporization Definition: Energy required to vaporize 1 mole of liquid Relationship to Intermolecular Forces: ​ Strong forces = Higher heat vaporization  ​ More energy needed to break molecular bonds Boiling Point Definition: Temperature where liquid converts to gas Characteristics: ​ Occurs when vapor pressure equals external pressure ​ Directly related to intermolecular force strength Comparative Analysis Table Property Weak IMF Substances Strong IMF Substances Vapor Pressure High Low Viscosity Low High Boiling Point Low HIgh Surface Tension Low High Intermolecular Forces Impact General Principles Stronger Intermolecular forces result in: ​ Higher surface tension ​ Greater viscosity ​ Lower vapor pressure ​ HIgher boiling point ​ HIgher heat vaporization PROPERTIES OF WATER (PART 4) Water: Unique Properties Solvent Capabilities Universal Solvent -​ Dissolves ionic and polar compounds -​ Can dissolve gases like oxygen and carbon dioxide Specific Heat High Specific Heat -​ Requires breaking hydrogen bonds to change temperature -​ Helps moderate Earth’s temperature Density Anomaly Unique Solid State -​ Less dense as solid than liquid -​ Hydrogen bonds create open structure -​ Causes ice to float on water Key Takeaway: Intermolecular forces fundamentally dictate liquid properties SOLIDS AND THEIR PROPERTIES (PART 5) Definition: Solid ​ A state of matter characterized by particles arranged in a fixed, ordered structure with a definite shape and volume. ​ Kinetic Molecular Theory: Describes solids in terms of: - Average kinetic energy - Distance among particles - Arrangement/order of particles - Attractive forces between particles Types of Solids 1.​ Crystalline Solids 2.​ Amorphous Solids Properties of Solids Crystalline Solids: Regular repeating 3D structure (crystal lattice) ​ Fixed geometric patterns or lattices ​ Examples: Ice, NaCl, diamond, graphite, sugar ​ Long-range order Amorphous Solids: No long-range order ​ Random orientation of particles ​ Examples: Glass, plastic, coal, rubber ​ Supercooled liquids with localized order Behavior When Heated​ Crystalline Solids ​ Specific melting point  ​ Sharp change in physical properties at melting point Amorphous Solids ​ Gradually soften when heated ​ Melt over a wide range of temperatures ​ Variation in particle arrangement causes uneven melting Types of Crystals (Crystalline Solids) ​ Metallic Crystals: ​ Atoms lose electrons to form positive ions (cations) ​ Held by metallic bonds (sea of electrons model) Ionic Crystals: ​ Made of cations and anions ​ Strong electrostatic interactions ​ High melting points Molecular Crystals: ​ Atoms or molecules held by hydrogen bonding/dipole-dipole and dispersion forces ​ Relatively low melting points Covalent Network Crystals: ​ Atoms covalently bonded to nearest neighbors ​ No individual molecules, entire crystals is one large molecule ​ Examples: Diamond (Cdiamond), Graphite (Cgraphite), SiO2, BN PHASE CHANGE AND PHASE DIAGRAM (PART 6) Phase Change​ Definition: Transformations of matter from one physical state to another Key Characteristics: ​ Occur when energy (usually heat) is added or removed ​ Characterized by changes in molecular order ​ Molecular Order Hierarchy: - Solid: Greatest order - Gas: Greatest randomness/disorder Types of Phase Changes Phase Transition Types Endothermic: Heat absorbed (heat gained) Exothermic: Heat released (heat lost) Specific Transitions 1.​ Melting 2.​ Vaporization 3.​ Sublimation 4.​ Condensation 5.​ Freezing 6.​ Deposition Energy and Phase Changes 1.​ Kinetic Energy Increase -​ Particles move faster -​ Temperature increases 2.​ Breaking Attractive Forces -​ No immediate temperature change -​ Physical appearance changes Heat Removal Effects 1.​ Kinetic Energy Decrease -​ Particle motion slows -​ Temperature decreases 2.​ Formation of Attractive Forces -​ Potential phase change -​ No temperature change Phase Diagram Definition: -​ Graphical representation of substance’’s physical states -​ Show combinations of pressure and temperature -​ Unique to each substance Diagram Components Three Primary Areas 1.​ Solid Region 2.​ Liquid Region 3.​ Vapor (Gas) Region Boundary Lines 1.​ Green Line: Solid-Liquid Transition -​ Represents melting and freezing points -​ Shows pressure’s effect on melting 2.​ Blue Line: Liquid-Gas Transition -​ Represents vaporization and condensation -​ Indicates pressure’s impact on boiling point 3.​ Red Line: Solid-Gas Transition -​ Represents sublimation and deposition Special Points 1.​ Triple Point -​ Unique pressure and temperature -​ All three matter phases coexist -​ Equilibrium point for solid-liquid-gas states 2.​ Critical Point -​ Liquid and gas phases merge -​ Creates supercritical fluid -​ Defined by critical temperature and pressure HEATING AND COOLING CURVE (PART 7) Heating Curve Characteristics ​ Temperature vs. Heat Added graph ​ Often uses time as x-axis ​ Reveals temperature and phase change behaviors Temperature Change Patterns 1.​ Temperature Changing Segments ○​ Indicates kinetic energy modifications ○​ Particles accelerating or decelerating 2.​ Constant Temperature Segments ○​ Phase changes occurring  ○​ Heat energy breaking/forming molecular bonds Phase Change Equilibrium ​ Two physical states coexist during transitions ​ Interconversion occurs at specific temperatures Recommended Study Strategies ​ Practice interpreting phase diagrams ​ Understand molecular behavior during transitions ​ Memorize key points and transition characteristics