General Chemistry Exam 3 Review: Atomic Structure, Periodic Properties, Chemical Bonding, and Molecular Geometry

Electronic Structure of Atoms (Chapter 6, part)

Quantum Theory and Atomic Structure

The quantum mechanical model explains the behavior of electrons in atoms, introducing concepts such as energy levels, orbitals, and electron configurations.

• Quantum Theory of the Atom: Describes electrons as wave-like entities with quantized energy levels.

• Energy Levels and Orbitals: Electrons occupy specific energy levels (shells) and subshells (orbitals: s, p, d, f).

• Quantum Numbers: Four quantum numbers (n, l, ml, ms) define the state of an electron.

• Electron Filling Order: Electrons fill orbitals in order of increasing energy (Aufbau principle).

• Condensed Electron Configurations: Shorthand notation using noble gas core to represent inner electrons.

Example: The electron configuration of sodium: [Ne] 3s1

Key Equation:

$E_n = -\frac{R_H}{n^2}$ (Energy of electron in hydrogen atom)

Periodic Properties of Elements (Chapter 7)

Effective Nuclear Charge and Atomic Trends

Periodic trends arise from the arrangement of electrons and the effective nuclear charge experienced by valence electrons.

• Effective Nuclear Charge (Zeff): The net positive charge experienced by an electron in a multi-electron atom.

• Atomic and Ionic Radii: Atomic radius decreases across a period and increases down a group; ionic radius depends on ion charge.

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

• Electron Affinity: The energy change when an atom gains an electron; generally becomes more negative across a period.

• Metallic Character: Metals lose electrons easily; metallic character increases down a group and decreases across a period.

• Physical and Chemical Properties: Trends in melting points, reactivity, and other properties are explained by periodic trends.

Example: Chlorine has a higher electron affinity than sodium.

Basic Concepts of Chemical Bonding (Chapter 8)

Types of Chemical Bonds and Lewis Structures

Chemical bonding involves the sharing or transfer of electrons between atoms to achieve stability.

• Ionic Bonds: Formed by transfer of electrons from metals to nonmetals, resulting in oppositely charged ions.

• Covalent Bonds: Formed by sharing of electrons between nonmetals.

• Polar vs. Nonpolar Covalent Bonds: Polar bonds have unequal sharing due to differences in electronegativity.

• Electronegativity: The ability of an atom to attract electrons in a bond; difference determines bond polarity.

• Lewis Structures: Diagrams showing bonding and lone pairs in molecules.

• Resonance: Some molecules have multiple valid Lewis structures; actual structure is a hybrid.

• Formal Charge: Used to determine the most stable Lewis structure.

• Bond Enthalpy: Energy required to break a bond; used to estimate reaction enthalpy.

Example: In water (H2O), oxygen is more electronegative than hydrogen, resulting in polar covalent bonds.

Key Equation:

$\Delta H_{rxn} = \sum \text{(bond enthalpies of bonds broken)} - \sum \text{(bond enthalpies of bonds formed)}$

Molecular Geometry and Bonding Theories (Chapter 9)

VSEPR Model and Hybridization

The shape of molecules is determined by the arrangement of electron pairs around the central atom, as explained by the VSEPR model and hybridization theory.

• VSEPR Model: Valence Shell Electron Pair Repulsion theory predicts molecular shapes based on repulsion between electron pairs.

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

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals (sp, sp2, sp3, etc.).

• Multiple Bonds: Sigma (σ) and pi (π) bonds formed by overlap of orbitals.

• Bond Angles: Determined by geometry; e.g., tetrahedral has 109.5° bond angles.

• Resonance and Delocalization: Some molecules have delocalized electrons, affecting geometry and stability.

Example: Methane (CH4) has a tetrahedral geometry due to sp3 hybridization.

HTML Table: Summary of Periodic Trends

Property	Trend Across Period	Trend Down Group
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electron Affinity	Becomes more negative	Becomes less negative
Metallic Character	Decreases	Increases

Additional info: Some details, such as specific examples and equations, were inferred to provide a complete study guide based on the review sheet topics.