Chemical Bonding, Molecular Geometry, and Bonding Theories: Study Guide

Chapter 8: Chemical Bonding

8.1: Lewis Dot Symbols

Lewis dot symbols are a simple way to represent the valence electrons of atoms. Each dot represents a valence electron, and the arrangement helps predict bonding behavior.

• Definition: A Lewis dot symbol consists of the element's symbol surrounded by dots representing valence electrons.

• Application: Used to visualize electron transfer in ionic bonding and electron sharing in covalent bonding.

• Example: The Lewis dot symbol for oxygen (O) is O with six dots around it.

8.2: Electronegativity

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond.

• Trends: Increases across a period (left to right) and decreases down a group in the periodic table.

• Pauling Scale: The most commonly used scale for electronegativity values.

• Example: Fluorine is the most electronegative element.

8.3: Ionic vs. Covalent Bonding

Chemical bonds can be classified as ionic or covalent based on how electrons are distributed between atoms.

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

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

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

• Nonpolar Covalent Bonds: Electrons are shared equally.

• Example: NaCl (ionic), H2O (polar covalent), Cl2 (nonpolar covalent).

8.4: Drawing Lewis Structures

Lewis structures represent the arrangement of electrons in molecules and polyatomic ions.

• Steps:

  1. Count total valence electrons.

  2. Arrange atoms with the least electronegative atom in the center (except hydrogen).

  3. Connect atoms with single bonds.

  4. Distribute remaining electrons to complete octets (or duets for hydrogen).

  5. Use double or triple bonds if necessary to satisfy the octet rule.

• Example: CO2 has two double bonds between carbon and each oxygen.

8.5: Formal Charge

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

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

• Application: The best Lewis structure has formal charges closest to zero and negative charges on the most electronegative atoms.

• Example: In the nitrate ion (NO3-), resonance structures distribute the negative charge.

8.6: Resonance

Some molecules cannot be represented by a single Lewis structure; instead, resonance structures are used.

• Definition: Resonance structures are two or more valid Lewis structures for the same molecule, differing only in the placement of electrons.

• Example: Ozone (O3) has two resonance structures.

8.7: Exceptions to the Octet Rule

Some molecules have atoms that do not follow the octet rule.

• Types of Exceptions:

  • Odd-electron species: Molecules with an odd number of electrons (e.g., NO).

  • Incomplete octet: Atoms with fewer than eight electrons (e.g., BF3).

  • Expanded octet: Atoms in period 3 or higher can have more than eight electrons (e.g., SF6).

Chapter 10: Molecular Geometry and Bonding Theories

10.1: Molecular Geometry (VSEPR Theory)

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

• Key Principle: Electron pairs (bonding and lone pairs) around a central atom arrange themselves to minimize repulsion.

• Common Geometries:

  • Linear: 180° bond angle (e.g., CO2)

  • Trigonal planar: 120° (e.g., BF3)

  • Tetrahedral: 109.5° (e.g., CH4)

  • Trigonal bipyramidal: 90°, 120° (e.g., PCl5)

  • Octahedral: 90° (e.g., SF6)

• Example: Ammonia (NH3) is trigonal pyramidal due to one lone pair on nitrogen.

10.2: Molecular Polarity and Dipole Moments

Molecular polarity depends on both the polarity of individual bonds and the geometry of the molecule.

• Dipole Moment: A measure of the separation of positive and negative charges in a molecule.

• Determining Polarity: A molecule is polar if it has polar bonds arranged asymmetrically, resulting in a net dipole moment.

• Example: Water (H2O) is polar; carbon dioxide (CO2) is nonpolar despite having polar bonds due to its linear geometry.

10.3: Valence Bond Theory and Hybridization

Valence bond theory explains how atomic orbitals combine to form chemical bonds. Hybridization describes the mixing of atomic orbitals to form new, equivalent hybrid orbitals.

• Types of Hybridization:

  • sp: Linear geometry

  • sp2: Trigonal planar geometry

  • sp3: Tetrahedral geometry

  • sp3d: Trigonal bipyramidal geometry

  • sp3d2: Octahedral geometry

• Sigma (σ) and Pi (π) Bonds: Single bonds are sigma bonds; double and triple bonds contain one sigma and one or two pi bonds, respectively.

• Example: Ethene (C2H4) has sp2 hybridization and a double bond (one σ and one π bond).

Table: Common Molecular Geometries and Hybridizations

Electron Domains	Geometry	Bond Angle	Hybridization	Example
2	Linear	180°	sp	CO2
3	Trigonal Planar	120°	sp2	BF3
4	Tetrahedral	109.5°	sp3	CH4
5	Trigonal Bipyramidal	90°, 120°	sp3d	PCl5
6	Octahedral	90°	sp3d2	SF6

Exam Preparation Tips

• Review how to draw Lewis structures and assign formal charges.

• Practice predicting molecular shapes using VSEPR theory.

• Understand the relationship between molecular geometry and polarity.

• Be able to identify hybridization and types of bonds in molecules.

• Refer to the provided textbook pages for detailed explanations and practice problems.