General Chemistry Study Notes: Chemical Bonding, Lewis Structures, and Molecular Geometry

Chemical Bonding and Lewis Structures

Lewis Dot Structures

Lewis dot structures are diagrams that represent the valence electrons of atoms within a molecule. They are used to visualize the arrangement of electrons and predict bonding between atoms.

• Valence Electrons: Electrons in the outermost shell, involved in chemical bonding.

• Lewis Symbols: Dots around an element's symbol represent its valence electrons.

• Example: Oxygen atom has 6 valence electrons, so its Lewis symbol is O with six dots around it.

Formal Charge

Formal charge is a bookkeeping tool used to determine the charge distribution within a molecule or ion. It helps identify the most stable Lewis structure.

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

• Application: Assign formal charges to atoms in polyatomic ions to find the most stable structure.

• Example: In the cyanide ion (CN-), formal charges help determine the best Lewis structure.

Ionic and Covalent Bonds

Chemical bonds are classified as ionic or covalent based on the nature of electron sharing or transfer between atoms.

• Ionic Bond: Formed by the transfer of electrons from a metal to a nonmetal, resulting in oppositely charged ions.

• Covalent Bond: Formed by the sharing of electrons between two nonmetals.

• Polar Covalent Bond: Unequal sharing of electrons, leading to partial charges.

• Example: NaCl is ionic, H2O is polar covalent.

Molecular Geometry and VSEPR Theory

VSEPR Theory

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on the repulsion between electron pairs around a central atom.

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

• Common Geometries:

  • Linear: 2 domains, 180° bond angle

  • Trigonal planar: 3 domains, 120° bond angle

  • Tetrahedral: 4 domains, 109.5° bond angle

  • Trigonal bipyramidal: 5 domains

  • Octahedral: 6 domains

• Example: CO2 is linear, NH3 is trigonal pyramidal.

Bond Types: Sigma and Pi Bonds

Covalent bonds can be classified as sigma (σ) or pi (π) bonds based on the type of orbital overlap.

• Sigma (σ) Bond: Formed by head-on overlap of orbitals; every single bond is a sigma bond.

• Pi (π) Bond: Formed by side-to-side overlap of p orbitals; present in double and triple bonds.

• Example: Ethylene (C2H4) has one sigma and one pi bond between the carbons.

Resonance Structures

Definition and Importance

Resonance structures are different Lewis structures for the same molecule that show the delocalization of electrons. The actual structure is a hybrid of all resonance forms.

• Resonance: Occurs when more than one valid Lewis structure can be drawn for a molecule.

• Delocalization: Electrons are spread over several atoms, increasing stability.

• Example: Ozone (O3) and nitrate ion (NO3-).

Bond Strength and Length

Comparing Bonds

Bond strength and length depend on the type of bond and the atoms involved.

• Single Bond: Longest and weakest

• Double Bond: Shorter and stronger than single bonds

• Triple Bond: Shortest and strongest

• Example: C≡C (triple bond) is stronger and shorter than C=C (double bond) or C–C (single bond).

Tables

Molecular Geometry Table

The following table summarizes molecular geometries based on the number of electron domains and lone pairs:

Electron Domains	Lone Pairs	Geometry
2	0	Linear
3	0	Trigonal planar
3	1	Bent
4	0	Tetrahedral
4	1	Trigonal pyramidal
4	2	Bent

Additional info:

• Some questions reference organic molecules and hybridization, which are covered in introductory general chemistry and organic chemistry.

• Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals suitable for bonding.

• For example, sp3 hybridization leads to tetrahedral geometry.