7.2 Covalent Bonding

Introduction to Chemical Bonds

Chemical bonds are the forces that hold atoms together in compounds. The two main types are ionic bonds and covalent bonds. Ionic bonds involve the transfer of electrons from one atom to another, typically between metals and nonmetals. Covalent bonds involve the sharing of electrons between nonmetal atoms.

• Ionic Bond: Electrostatic attraction between charged species due to electron transfer.

• Covalent Bond: Sharing of electrons between atoms.

• Electronegativity: Measure of an atom's ability to attract electrons in a covalent bond.

Bond Polarity: Determined by differences in electronegativity.

- : Nonpolar covalent bond

- : Polar covalent bond

- : Ionic bond

Example: NaCl (ionic), Cl2 (nonpolar covalent), HCl (polar covalent)

7.3 Dipole Moments and Molecular Shape

Bond Polarity and Dipole Moments

A dipole moment arises in molecules with polar covalent bonds and an asymmetric shape. The dipole moment () is a measure of the separation of positive and negative charges in a molecule.

• Dipole Moment Equation: where is the magnitude of the charge and is the distance between charges.

• To be polar, a molecule must:

  1. Contain at least one polar covalent bond.

  2. Have a molecular structure such that the sum of the vectors of each bond dipole moment does not cancel.

Examples: H2O (polar), CO2 (nonpolar), NH3 (polar)

7.4 Lewis Symbols and Structures

Lewis Symbols and Lewis Structures

Lewis symbols represent valence electrons as dots around the element symbol. Lewis structures show how atoms are bonded and the arrangement of valence electrons.

• Steps to Draw Lewis Structures:

  1. Write the element symbols and count valence electrons.

  2. Connect atoms with single bonds (pairs of electrons).

  3. Distribute remaining electrons to complete octets (or duets for H).

  4. Use double/triple bonds if necessary to satisfy octet rule.

  5. Check for formal charges and minimize them.

• Helpful Hints:

  • Carbon: 4 bonds (no lone pairs)

  • Nitrogen: 3 bonds, 1 lone pair

  • Oxygen: 2 bonds, 2 lone pairs

  • Hydrogen: 1 bond, no lone pairs

Examples: NH3, CH4, H2O, CO2

Exceptions to the Octet Rule

• Elements with fewer than 8 electrons (e.g., H, B, Be)

• Expanded octets (elements in period 3 or higher)

• Odd-electron molecules (radicals)

7.4 Formal Charges and Resonance

Formal Charges

Formal charge is a bookkeeping method to determine the charge distribution in a molecule.

• Formal Charge Equation:

• Structures with formal charges closest to zero are preferred.

• No adjacent positive and negative formal charges.

Resonance Structures

Resonance occurs when more than one valid Lewis structure can be drawn for a molecule. The actual structure is a hybrid of all resonance forms.

7.5 Strength of Ionic and Covalent Bonds

Bond Energy and Enthalpy

Bond energy is the energy required to break a bond. Bond enthalpy is used to estimate the enthalpy change of reactions.

• Bond Enthalpy Equation:

• Average bond energies can be used to estimate reaction enthalpies.

Example: Calculating for CH4 + 2Cl2 → CH2Cl2 + 2HCl using bond energies.

7.6 Molecular Structure and Polarity

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory predicts the 3D arrangement of atoms in a molecule by minimizing electron pair repulsion around the central atom.

• Basic electron group geometries:

  • Linear

  • Trigonal planar

  • Tetrahedral

  • Trigonal bipyramidal

  • Octahedral

• Bond angles depend on the number of electron pairs and lone pairs.

Electron-Pair Geometry vs. Molecular Structure

• Electron-pair geometry: Arrangement of all electron groups (bonding and lone pairs).

• Molecular structure: Arrangement of only atoms (ignoring lone pairs).

Example: H2O has tetrahedral electron-pair geometry but bent molecular structure.

Summary Tables

Summary of Molecular Geometries

The following tables summarize the relationship between steric number, number of ligands, lone pairs, electron group geometry, molecular shape, bond angles, dipole moment, and examples.

Summary of Molecular Geometries, continued