Study Notes: Chemical Bonding (Chapter 10)

Chemical Bonding

Electron Dot Formulas

Chemical bonding involves the sharing or transfer of electrons between atoms to achieve stable electron configurations. Electron dot formulas (also called Lewis structures) are a visual representation of the valence electrons in molecules and ions.

• Valence Electrons: Electrons in the outermost shell of an atom, involved in bonding.

• Lewis Dot Symbols: Dots are placed around the chemical symbol to represent valence electrons.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (an octet), similar to noble gases.

• Duet Rule: Hydrogen and helium are stable with two electrons.

Steps for Drawing Lewis Structures:

1. Count the total number of valence electrons for all atoms in the molecule or ion.

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

3. Connect atoms with single bonds (each bond = 2 electrons).

4. Distribute remaining electrons as lone pairs to complete octets (or duets for H).

5. If necessary, form double or triple bonds to satisfy the octet rule.

Example: Draw the electron dot formulas for the following molecules: F2, CH4, N2, O2, CO2, H2O, CH3Cl, NO3-, SO2 (resonance).

Predicting the Shapes of Molecules (VSEPR Theory)

The shape of a molecule is determined by the repulsion between electron pairs (bonding and lone pairs) around the central atom. This is described by the Valence Shell Electron Pair Repulsion (VSEPR) theory.

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

• VSEPR Principle: Electron domains arrange themselves as far apart as possible to minimize repulsion.

• Common Molecular Shapes:

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

  • Trigonal planar: 120° bond angle (e.g., SO2)

  • Tetrahedral: 109.5° bond angle (e.g., CH4)

  • Trigonal pyramidal: ~107° bond angle (e.g., NH3)

  • Bent: <120° or <109.5° bond angle (e.g., H2O)

Example: Give the shape of the molecules and ions in the previous example.

Polar Covalent Bonds

Covalent bonds involve the sharing of electrons between atoms. If the sharing is unequal due to differences in electronegativity, the bond is polar.

• Electronegativity: The ability of an atom to attract shared electrons in a bond.

• Polar Bond: A bond in which electrons are shared unequally, resulting in partial positive (δ+) and partial negative (δ-) charges.

• Nonpolar Bond: A bond in which electrons are shared equally.

• Dipole: A molecule with a positive end and a negative end due to uneven electron distribution.

Example: Determine whether the following bonds are polar or nonpolar: H–Cl, F–F, Na–Cl, H–F, C–H.

Molecular Polarity

The polarity of a molecule depends on both the polarity of its bonds and its shape.

• If the molecule is symmetrical, dipoles may cancel, resulting in a nonpolar molecule (e.g., CO2).

• If the molecule is asymmetrical, dipoles do not cancel, resulting in a polar molecule (e.g., H2O).

Example: For each molecule, determine the shape and polarity: CH3Cl, H2O, PCl3, ClO3-, ClO2.

Summary Table: Molecular Geometry and Polarity

Molecule/Ion	Lewis Structure	Shape	Polarity
CO2	O=C=O	Linear	Nonpolar
H2O	H–O–H (with 2 lone pairs on O)	Bent	Polar
CH4	Central C with 4 H atoms	Tetrahedral	Nonpolar
NH3	Central N with 3 H atoms and 1 lone pair	Trigonal pyramidal	Polar
CH3Cl	Central C with 3 H and 1 Cl	Tetrahedral	Polar
ClO3-	Central Cl with 3 O atoms and 1 lone pair	Trigonal pyramidal	Polar
ClO2	Central Cl with 2 O atoms and 2 lone pairs	Bent	Polar

Key Formulas and Concepts

• Octet Rule: Atoms gain, lose, or share electrons to achieve 8 valence electrons.

• VSEPR Theory: Electron pairs arrange to minimize repulsion, determining molecular shape.

• Electronegativity Difference: Determines bond polarity.

  • Nonpolar covalent: ΔEN < 0.5

  • Polar covalent: 0.5 ≤ ΔEN < 2.0

  • Ionic: ΔEN ≥ 2.0

Sample Lewis Structures and Shapes

• CO2: $O=C=O$ (linear, nonpolar)

• H2O: $H-O-H$ with two lone pairs on O (bent, polar)

• NH3: $\mathrm{H-N-H}$ with one lone pair on N (trigonal pyramidal, polar)

• CH4: Central C with four H atoms (tetrahedral, nonpolar)

Practice Problems

1. Draw Lewis structures for the following: F2, CH4, N2, O2, CO2, H2O, CH3Cl, NO3-, SO2 (resonance).

2. Predict the shape and polarity of each molecule using VSEPR theory.

3. Classify each bond as polar or nonpolar and indicate the direction of the dipole if polar.

Additional info: These notes expand on the provided class concepts with definitions, examples, and a summary table for clarity and exam preparation.