General Chemistry II Exam III Review: Chemical Bonding, Intermolecular Forces, and Solutions

Chapter 10: Chemical Bonding I – The Lewis Model

Bond Length and Bond Energy

Bond length is the average distance between the nuclei of two bonded atoms. Bond energy (or bond enthalpy) is the energy required to break one mole of a particular bond in a gaseous substance.

• Bond Length: Shorter bonds are generally stronger; triple bonds < double bonds < single bonds.

• Bond Energy: Higher bond energy means a stronger bond.

• Example: The bond energy of an O-H bond in water is about 463 kJ/mol.

Enthalpy Changes in Reactions (Bond Enthalpies)

Bond enthalpy can be used to estimate the enthalpy change ($\Delta H$) of a reaction by considering the bonds broken and formed.

• Formula: $\Delta H = \sum \text{(bond energies of bonds broken)} - \sum \text{(bond energies of bonds formed)}$

• Application: Used to estimate reaction energetics when standard enthalpies of formation are unavailable.

Coulomb's Law and Lattice Energy

Coulomb's Law describes the electrostatic interaction between charged particles. Lattice energy is the energy released when ions form a crystalline lattice.

• Formula: $E = \dfrac{kQ_1Q_2}{r}$

• Lattice Energy: Higher for ions with greater charge and smaller ionic radii.

Lewis Structures and Resonance

Lewis structures represent the arrangement of electrons in molecules. Resonance occurs when more than one valid Lewis structure can be drawn for a molecule.

• Resonance: Delocalization of electrons increases stability.

• Example: Ozone (O3) has two resonance structures.

Exceptions to the Octet Rule

Some molecules do not follow the octet rule due to odd numbers of electrons, expanded octets, or incomplete octets.

• Expanded Octet: Elements in period 3 or higher can have more than 8 electrons (e.g., SF6).

• Incomplete Octet: Molecules like BF3 have less than 8 electrons around the central atom.

Chapter 11: Chemical Bonding II – Molecular Shapes, VSEPR & MO Theory

VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the geometry of molecules based on electron pair repulsion.

• Key Shapes: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Example: CH4 is tetrahedral.

Electron and Molecular Geometry

Electron geometry considers all electron groups; molecular geometry considers only atoms.

• Electron Geometry: Arrangement of all electron groups.

• Molecular Geometry: Arrangement of atoms only.

Polarity of Molecules

Molecular polarity depends on bond polarity and molecular shape.

• Polar Molecules: Have a net dipole moment (e.g., H2O).

• Nonpolar Molecules: Symmetrical molecules with no net dipole (e.g., CO2).

Hybridization and Molecular Orbitals

Atomic orbitals combine to form hybrid orbitals (sp, sp2, sp3) and molecular orbitals (bonding and antibonding).

• Hybridization: Explains molecular shapes and bonding.

• Molecular Orbital Theory: Describes electrons in molecules as delocalized over the entire molecule.

Chapter 12: Liquids, Solids & Intermolecular Forces

Properties of States of Matter

Solids, liquids, and gases differ in particle arrangement, energy, and intermolecular forces.

• Solids: Fixed shape and volume; strong intermolecular forces.

• Liquids: Fixed volume, variable shape; moderate forces.

• Gases: Variable shape and volume; weak forces.

Intermolecular Forces

Intermolecular forces determine physical properties like boiling point and solubility.

• Types: Dispersion (London), dipole-dipole, hydrogen bonding, ion-dipole.

• Hydrogen Bonding: Strongest type, occurs in molecules with N-H, O-H, or F-H bonds.

Phase Changes and Diagrams

Phase changes include melting, freezing, vaporization, condensation, and sublimation. Phase diagrams show the state of a substance at various temperatures and pressures.

• Heating Curve: Shows temperature change as heat is added.

• Phase Diagram: Indicates regions of solid, liquid, and gas.

Vaporization and Vapor Pressure

Vaporization is the process of a liquid becoming a gas. Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid.

• Clausius-Clapeyron Equation: Relates vapor pressure and temperature.

• Formula: $\ln P = -\dfrac{\Delta H_{vap}}{R} \left( \dfrac{1}{T} \right) + C$

Chapter 14: Solutions

Types and Properties of Solutions

Solutions are homogeneous mixtures of solute and solvent. Types include solid, liquid, and gaseous solutions.

• Solute: Substance dissolved.

• Solvent: Substance doing the dissolving.

• Example: Saltwater (NaCl in H2O).

Solubility and Saturation

Solubility is the maximum amount of solute that can dissolve in a solvent at a given temperature. Saturated solutions contain the maximum amount; unsaturated contain less; supersaturated contain more than the maximum.

• Factors Affecting Solubility: Temperature, pressure, nature of solute and solvent.

Concentration Units

Concentration expresses the amount of solute in a given amount of solution.

• Molarity (M): $M = \dfrac{\text{moles of solute}}{\text{liters of solution}}$

• Molality (m): $m = \dfrac{\text{moles of solute}}{\text{kg of solvent}}$

• Mass Percent: $\% = \dfrac{\text{mass of solute}}{\text{mass of solution}} \times 100$

Colligative Properties

Colligative properties depend on the number of solute particles, not their identity. These include boiling point elevation, freezing point depression, vapor pressure lowering, and osmotic pressure.

• Boiling Point Elevation: $\Delta T_b = iK_b m$

• Freezing Point Depression: $\Delta T_f = iK_f m$

• Raoult's Law: $P_{solution} = X_{solvent} P^0_{solvent}$

Preparing Solutions and Titrations

Preparing solutions involves dissolving a known amount of solute in solvent. Titrations are used to determine the concentration of an unknown solution.

• Titration: Addition of a solution of known concentration to react with a solution of unknown concentration.

Summary Table: Types of Intermolecular Forces

Type	Description	Example
Dispersion (London)	Temporary dipoles in all molecules	He, N2
Dipole-Dipole	Permanent dipoles in polar molecules	HCl, SO2
Hydrogen Bonding	H bonded to N, O, or F	H2O, NH3
Ion-Dipole	Ions interacting with polar molecules	Na+ in H2O

Additional info: These notes expand on the listed review topics by providing definitions, formulas, and examples for each concept, ensuring a comprehensive study guide for General Chemistry II exam preparation.