General Chemistry Study Guide: Electronic Structure, Periodic Properties, Chemical Bonding, and Molecular Geometry

Electronic Structure of Atoms

Electron Filling Order and Quantum Numbers

The arrangement of electrons in atoms follows specific rules based on energy levels and quantum numbers. Understanding these principles is essential for predicting chemical behavior.

• Aufbau Principle: Electrons fill orbitals starting from the lowest energy level to higher ones. The order is generally: $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \rightarrow 4p \rightarrow 5s$.

• Electron Configuration: The notation shows the distribution of electrons among orbitals. For example, germanium (Ge) has the configuration.

• Quantum Numbers: Four quantum numbers (n, l, ml, ms) describe the properties of electrons in atoms. Allowed values are:

  • Principal quantum number (n): positive integers (1, 2, 3...)

  • Angular momentum quantum number (l): integers from 0 to n-1

  • Magnetic quantum number (ml): integers from -l to +l

  • Spin quantum number (ms): +1/2 or -1/2

• Maximum Electrons per Energy Level: The maximum number of electrons in a shell is.

Example: The electron configuration for is.

Periodic Properties of the Elements

Atomic and Ionic Radii, Ionization Energy, and Electronegativity

Periodic trends help predict the physical and chemical properties of elements.

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

• Ionic Radius: Cations are smaller than their parent atoms; anions are larger.

• Ionization Energy: The energy required to remove an electron from an atom. Increases across a period and decreases within a group.

• Electronegativity: A measure of an atom's ability to attract electrons in a bond. Fluorine is the most electronegative element.

Example: The order of decreasing first ionization energy for C, N, O, Na is: N > O > C > Na.

Electron Configurations of Ions

When atoms form ions, electrons are added or removed according to the Aufbau principle and Hund's rule.

• Cations: Electrons are removed first from the highest principal quantum number shell.

• Anions: Electrons are added to the lowest available energy orbital.

Example: The electron configuration for is.

Basic Concepts of Chemical Bonding

Types of Chemical Bonds and Lewis Structures

Chemical bonds form when atoms share or transfer electrons to achieve stable electron configurations.

• Ionic Bonds: Formed by the transfer of electrons from a metal to a nonmetal, resulting in cations and anions.

• Covalent Bonds: Formed by sharing electrons between nonmetals.

• Polar Covalent Bonds: Electrons are shared unequally due to differences in electronegativity.

• Lewis Structures: Diagrams that show the arrangement of valence electrons among atoms in a molecule.

Example: The Lewis structure for HCN is: H–C≡N

Bond Polarity and Electronegativity

The difference in electronegativity between two atoms determines the bond type.

• Nonpolar Covalent: Electronegativity difference < 0.5

• Polar Covalent: Electronegativity difference between 0.5 and 1.7

• Ionic: Electronegativity difference > 1.7

Example: If element A has an electronegativity 3.0 and B has 2.1, AB is polar covalent, with A partially negative and B partially positive.

Molecular Geometry and Bonding Theories

VSEPR Theory and Molecular Shapes

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

• Electron Domains: Regions of electron density (bonds or lone pairs) around a central atom.

• Common Geometries:

  • 2 domains: Linear (180°)

  • 3 domains: Trigonal planar (120°)

  • 4 domains: Tetrahedral (109.5°)

  • 5 domains: Trigonal bipyramidal (90°, 120°)

  • 6 domains: Octahedral (90°)

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals. For example, sp2 hybridization leads to trigonal planar geometry.

Example: The geometry of a sp3-hybridized atom is tetrahedral.

Formal Charge and Resonance

Formal charge helps determine the most stable Lewis structure for a molecule.

• Formal Charge Formula: $\text{Formal charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2} \times \text{Bonding electrons}$

Example: In HCN, the formal charge on C is 0.

Bond Dissociation Energy

Bond dissociation energy is the energy required to break a bond in a molecule. It can be used to estimate reaction enthalpies.

• Calculation: $\Delta H = \sum \text{Bond energies of reactants} - \sum \text{Bond energies of products}$

Example: Using bond energies, calculate for a reaction by summing the energies of bonds broken and subtracting the energies of bonds formed.

Properties of Solutions and Chemical Bonding

Expanded Valence Shells and Exceptions to the Octet Rule

Some elements can have more than eight electrons in their valence shell, especially those in period 3 and beyond.

• Expanded Octet: Elements like P, S, and Cl can form compounds with more than eight valence electrons.

• Octet Rule Exceptions: Molecules with odd numbers of electrons, incomplete octets, or expanded octets.

Example: $SF_6$ has 12 electrons around sulfur.

Summary Table: Periodic Trends

Property	Across a Period	Down a Group
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electronegativity	Increases	Decreases
Metallic Character	Decreases	Increases

Additional info:

• Some questions reference specific elements and ions; students should consult a periodic table for atomic numbers and electron configurations.

• Hybridization and molecular geometry are covered in detail in chapters on bonding theories and molecular shapes.

• Bond dissociation energy calculations require knowledge of bond energies, typically provided in tables.