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 gaseous molecules.

• 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

Enthalpy changes can be estimated using bond energies. The enthalpy change for a reaction is the difference between the energy required to break bonds in reactants and the energy released when new bonds form in products.

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

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

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 molecular shapes based on electron pair repulsion around the central atom.

• Electron Groups: Lone pairs and bonds count as electron groups.

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

• Example: CH4 is tetrahedral.

Polarity of Molecules

Molecular polarity depends on bond polarity and molecular geometry.

• Polar Molecules: Have an uneven distribution of charge (e.g., H2O).

• Nonpolar Molecules: Symmetrical molecules with even charge distribution (e.g., CO2).

Molecular Orbital Theory

Molecular orbital (MO) theory describes electrons in molecules as occupying orbitals that extend over the entire molecule.

• Bonding and Antibonding Orbitals: Bonding orbitals stabilize molecules; antibonding orbitals destabilize.

• Application: Explains magnetic properties and bond order.

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 (IMFs) are attractions between molecules, affecting boiling/melting points and solubility.

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

• Hydrogen Bonding: Strong IMF in molecules with N-H, O-H, or F-H bonds.

• Example: Water's high boiling point is due to hydrogen bonding.

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.

• Boiling Point: Temperature at which vapor pressure equals external pressure.

• Clausius-Clapeyron Equation: Relates vapor pressure and temperature: $\ln P = -\frac{\Delta H_{vap}}{R} \left( \frac{1}{T} \right) + C$

• Heating/Cooling Curves: Show temperature changes during phase transitions.

Chapter 14: Solutions

Types and Properties of Solutions

Solutions are homogeneous mixtures of solute and solvent. Properties depend on concentration and nature of components.

• Types: Saturated, unsaturated, supersaturated.

• Concentration Units: Molarity ($M$), molality ($m$), mole fraction ($X$).

• Example: A 1 M NaCl solution contains 1 mole of NaCl per liter of solution.

Solubility and Factors Affecting Solubility

Solubility is the maximum amount of solute that can dissolve in a solvent at a given temperature.

• Factors: Temperature, pressure (for gases), nature of solute/solvent.

• Henry's Law: $C = kP$ (solubility of a gas is proportional to its pressure).

Colligative Properties

Colligative properties depend on the number of solute particles, not their identity.

• Boiling Point Elevation: $\Delta T_b = i K_b m$

• Freezing Point Depression: $\Delta T_f = i K_f m$

• Vapor Pressure Lowering: Addition of solute lowers vapor pressure.

• 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.

• Mole Fraction: $X_A = \frac{n_A}{n_{total}}$

• Titration: Involves adding a titrant to a solution until the reaction is complete (endpoint).

Summary Table: Types of Intermolecular Forces

Type of Force	Example	Relative Strength
Dispersion (London)	All molecules (e.g., N2)	Weakest
Dipole-Dipole	Polar molecules (e.g., HCl)	Intermediate
Hydrogen Bonding	H2O, NH3	Strongest (among IMFs)
Ion-Dipole	Na+ in water	Very strong (in solutions)

Additional info: These notes expand on the listed review topics by providing definitions, formulas, and examples for each major concept. For full mastery, students should practice drawing Lewis structures, calculating colligative properties, and interpreting phase diagrams.