General Chemistry II: Study Guide for Chapters 11–19 (Intermolecular Forces, Solids, Solutions, Kinetics, Equilibria, Thermodynamics, and Environmental Chemistry)

Chapter 11: Liquids and Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces are the forces of attraction between molecules, which influence the physical properties of substances such as boiling and melting points.

• London Dispersion Forces: Present in all molecules, but are the only forces in nonpolar molecules.

• Dipole-Dipole Forces: Occur between polar molecules.

• Hydrogen Bonding: A special, strong dipole-dipole interaction involving H bonded to N, O, or F.

Example: CH4 exhibits only London dispersion forces, while H2O exhibits hydrogen bonding, dipole-dipole, and dispersion forces.

Properties of Liquids

• As a gaseous element condenses, atoms become closer and have stronger attractions for one another.

• Hydrogen bonding is the predominant intermolecular force in molecules like H2O and NH3.

Ion-Dipole Forces

When ionic compounds dissolve in water, the attraction between ions and polar water molecules is called ion-dipole interaction.

Phase Changes and Heating Curves

• During phase changes, energy is used to increase distances between molecules (overcoming intermolecular forces).

• The slope of a segment in a heating curve corresponds to the heat capacity of the phase present.

Phase Diagrams

• Show the state of a substance at various temperatures and pressures.

• The normal boiling point is the temperature at which the vapor pressure equals 1 atm.

Chapter 12: Solids and Modern Materials

Types of Solids

• Crystalline Solids: Have long-range order and repeating patterns.

• Amorphous Solids: Lack long-range order.

Unit Cells and Packing Efficiency

• Unit Cell: The smallest repeating unit in a crystal lattice.

• Packing Efficiency: Fraction of volume in a crystal actually occupied by atoms.

Types of Bonding in Solids

• Ionic Solids: Held together by electrostatic attraction between cations and anions.

• Molecular Solids: Held together by intermolecular forces.

• Metallic Solids: Metal atoms held together by a 'sea' of delocalized electrons.

• Covalent-Network Solids: Atoms held together by covalent bonds in a network.

Semiconductors

• n-type: Doping with atoms that have more valence electrons than the host.

• p-type: Doping with atoms that have fewer valence electrons than the host.

Chapter 13: Properties of Solutions

Solution Formation

• Solutions form when intermolecular forces between solute and solvent are comparable to those within each substance.

• Saturated Solution: Contains the maximum amount of solute at a given temperature.

• Supersaturated Solution: Contains more solute than is present in a saturated solution.

Solubility and Henry's Law

• "Like dissolves like": Polar solutes dissolve in polar solvents, nonpolar in nonpolar.

• Henry's Law: $C = k_P P$ (solubility is proportional to partial pressure of gas above solution).

Concentration Units

• Molarity (M): Moles of solute per liter of solution.

• Molality (m): Moles of solute per kilogram of solvent.

• Mass Percent, Mole Fraction, and Molarity are used to express solution concentration.

Colligative Properties

• Depend on the number of solute particles, not their identity (e.g., boiling point elevation, freezing point depression).

• Freezing Point Depression: $riangle T_f = K_f m$

Chapter 14: Chemical Kinetics

Reaction Rates

• Rate of reaction is the change in concentration of reactant or product per unit time.

• Rate Law: $ext{Rate} = k[ ext{A}]^m[ ext{B}]^n$

• Order of reaction is determined experimentally.

Half-Life

• Time required for half of a reactant to be consumed.

• For first-order reactions: $t_{1/2} = rac{0.693}{k}$

Reaction Mechanisms

• Sequence of elementary steps by which a reaction occurs.

• The slowest step determines the rate law (rate-determining step).

Activation Energy

• The minimum energy required for a reaction to occur.

Chapter 16: Acid-Base Equilibria

Brønsted-Lowry Acids and Bases

• Acid: Proton (H+) donor.

• Base: Proton (H+) acceptor.

Conjugate Acid-Base Pairs

• Acids and bases exist in pairs that differ by one proton.

pH and pOH

• $ext{pH} = - ext{log}[ ext{H}^+]$

• $ext{pOH} = - ext{log}[ ext{OH}^-]$

• $ext{pH} + ext{pOH} = 14$ at 25°C

Acid and Base Strength

• Strong acids/bases dissociate completely in water; weak acids/bases only partially dissociate.

• $K_a$ and $K_b$ are the acid and base dissociation constants, respectively.

Chapter 17: Additional Aspects of Aqueous Equilibria

Buffer Solutions

• Mixtures of a weak acid and its conjugate base (or weak base and conjugate acid) that resist changes in pH.

• Henderson-Hasselbalch Equation: $ext{pH} = ext{p}K_a + ext{log} rac{[ ext{A}^-]}{[ ext{HA}]}$

Solubility Equilibria

• The solubility product constant ($K_{sp}$) describes the equilibrium between a solid and its ions in solution.

Chapter 18: Chemistry of the Environment

Atmospheric Chemistry

• CO2 from fossil fuels contributes to the greenhouse effect and global warming.

• Ozone (O3) in the stratosphere protects Earth from harmful UV radiation.

• Photodissociation and reactions involving nitrogen and oxygen species affect atmospheric composition.

Chapter 19: Chemical Thermodynamics

Spontaneity and Entropy

• A spontaneous process occurs without outside intervention.

• Entropy (S): A measure of disorder or randomness.

• Second Law of Thermodynamics: The entropy of the universe increases in a spontaneous process.

Gibbs Free Energy

• $riangle G = riangle H - T riangle S$

• If $riangle G < 0$, the process is spontaneous.

Standard Free Energy Changes

• Calculated using standard enthalpy and entropy values.

Equilibrium and Thermodynamics

• Relationship between $riangle G$ and the equilibrium constant $K$: $riangle G^ ext{°} = -RT ext{ln} K$

Tables

Sample Table: Thermodynamic Quantities for Selected Substances at 298.15 K (25°C)

Substance	ΔHf° (kJ/mol)	S° (J/mol·K)	ΔGf° (kJ/mol)
CO2(g)	-393.5	213.6	-394.4
H2O(l)	-285.8	69.9	-237.1
SO2(g)	-296.8	248.2	-300.4
SO3(g)	-395.7	256.8	-370.4
S(s, rhombic)	0	31.8	0

Additional info: Table values are used to calculate ΔG, ΔH, and ΔS for reactions involving these substances.