General Chemistry: Atomic Structure, Chemical Bonding, Gases, and Intermolecular Forces (Ch. 9–12 Study Guide)

Chapter 9: Electrons in Atoms and The Periodic Table

Introduction

This chapter explores the electromagnetic radiation associated with the electronic structure of atoms. It covers electron configurations, orbital diagrams, and periodic trends such as ionization energy, atomic size, and metallic character.

Key Concepts

• Electromagnetic Radiation: Energy that travels through space as waves; includes visible light, ultraviolet, infrared, etc.

• Wavelength, Frequency, and Energy: Related by the equation $c = \lambda \nu$, where $c$ is the speed of light, $\lambda$ is wavelength, and $\nu$ is frequency.

• Electron Configuration: The arrangement of electrons in an atom's orbitals, e.g., $1s^2 2s^2 2p^6$ for neon.

• Orbital Diagrams: Visual representations showing electron spins in orbitals.

• Valence vs. Core Electrons: Valence electrons are in the outermost shell and participate in chemical bonding; core electrons are closer to the nucleus.

• Periodic Trends:

  • Atomic Size: Generally increases down a group and decreases across a period.

  • Ionization Energy: Energy required to remove an electron; increases across a period, decreases down a group.

  • Metallic Character: Tendency to lose electrons; increases down a group, decreases across a period.

Example

• Predicting atomic size: Sodium ($Na$) is larger than chlorine ($Cl$) because it is further left in the period.

Chapter 10: Chemical Bonding

Introduction

This chapter examines Lewis structures for atoms and simple molecules, including concepts of polarity and resonance. It discusses how electron pair repulsion theory predicts molecular shape and polarity.

Key Concepts

• Lewis Structures: Diagrams showing the bonding between atoms and lone pairs of electrons.

• Writing Lewis Structures: For elements, ionic compounds, and covalent compounds.

• Resonance Structures: Multiple valid Lewis structures for a molecule, e.g., $O_3$ (ozone).

• VSEPR Theory: Valence Shell Electron Pair Repulsion theory predicts molecular shapes based on electron pair repulsion.

• Polarity: Determined by differences in electronegativity and molecular geometry.

Example

• Water ($H_2O$) is polar due to its bent shape and difference in electronegativity between $H$ and $O$.

Chapter 11: Gases

Introduction

This chapter covers kinetic molecular theory and the properties of gases. It explains relationships among pressure, volume, temperature, and amount of gas, and introduces gas laws and partial pressures.

Key Concepts

• Pressure Units: Atmospheres (atm), Pascals (Pa), mmHg, torr.

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant $T$ and $n$)

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant $P$ and $n$)

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Ideal Gas Law: $PV = nRT$

• Dalton's Law of Partial Pressures: $P_{total} = P_1 + P_2 + ... + P_n$

• Stoichiometric Calculations: Using gas laws to relate moles, volume, and pressure in chemical reactions.

Example

• Calculating the volume of $O_2$ produced at STP from the decomposition of $KClO_3$.

Chapter 12: Liquids, Solids, and Intermolecular Forces

Introduction

This chapter focuses on interactions between molecules, including intermolecular forces and their influence on the properties of liquids and solids. It covers phase changes and related calculations.

Key Concepts

• Intermolecular Forces:

  • London Dispersion Forces: Present in all molecules, especially nonpolar.

  • Dipole-Dipole Forces: Occur between polar molecules.

  • Hydrogen Bonding: Strong dipole-dipole interaction involving $H$ bonded to $N$, $O$, or $F$.

• Phase Changes: Melting, boiling, vaporization, fusion.

• Heat of Vaporization ($\Delta H_{vap}$): Energy required to vaporize a liquid.

• Heat of Fusion ($\Delta H_{fus}$): Energy required to melt a solid.

• Calculations: Use $q = mC\Delta T$ for temperature changes and $q = n\Delta H$ for phase changes.

• Physical Properties: Intermolecular forces affect melting and boiling points.

Example

• Water has a high boiling point due to strong hydrogen bonding.